Power point presentation for General Chemistry

1. General
Chemistry
1
ALYSSA MAE ANTONIO, LPT.

2. 1. Use quantum numbers to describe an
electron in an atom. (STEM_GC11ESIIa-b-54)
2. Determine the magnetic property of the
atom based on its electronic configuration.
(STEM_GC11ESIIa-b-57)
3. Draw an orbital diagram to represent the
electronic configuration of atoms.
(STEM_GC11ESIIa-b-58) 2
Learning Competencies:

3. Quantum
Numbers

4. - Quantum numbers describe values of
conserved quantities of a quantum system.
- An important aspect of quantum mechanics
is the quantization of observable quantities,
which is possible through quantum
numbers since they are discrete sets of
integers and half-integers (they can
approach infinity in some cases).
- This makes quantum mechanics different
from classical mechanics where mass,
charge, or momentum values range
4

5. - Particles have quantum numbers that can
represent their properties.
- This is because the set of quantum
numbers is unique for every electron.
- Pauli's exclusion principle states that no
two identical electrons may occupy the
same quantum state in an atom
simultaneously.
- Thus, no two electrons in an atom can have
exactly the same set of quantum numbers.
5

6. 6

7. 7
The Four types of
quantum numbers:

8. - This describes the energy level of an
electron.
- The value of n ranges from 1 to the electron
shell or energy level containing the
outermost electron of the atom.
- The larger the value of n, the farther the
distance of the electron in the orbital from
the nucleus. 8
1. Principal quantum
number (n)

9. - Also, as the distance from the nucleus
increases, the energy of each orbital
increases.
- This means that higher n values have
higher energies.
9
1. Principal quantum
number (n)

10. Example: Fluorine (F)
- Fluorine has an electron configuration of
1s2
2s2
2p5
.
- Through its electron configuration, we can tell
that its outermost electrons or valence
electrons are in the second energy level.
- Therefore, an electron in fluorine can have an
n value of 1 to 2.

11. - Angular quantum number is also called
azimuthal quantum number.
- This describes the shape of the orbital and
the magnitude of the orbital angular
momentum.
- l has values ranging from 0 to (n - 1).
11
2. Angular quantum
number (ℓ)

12. - Thus, the value of l depends on the value of
n. If n = 1, there is one possible value of l,
and that is 0. If n = 2, there are two possible
values of l, which are 0 and 1. If n = 3, the
three possible values of l are 0, 1, and 2.
- The values of I can also be determined by
the designated orbital of its valence
electrons. 12
2. Angular quantum
number (ℓ)

14. Example: Fluorine (5F)
- Since fluorine has an n value of 2, its possible
{ values are only 0 and 1.
- From its electron configuration, it has its
valence electrons in the p orbital.
- Therefore, its angular quantum number is 1.

15. - This describes the orientation of the
electron cloud.
- The value of m, depends on the value of l.
Given a value l, the range of m, values is
from -l to +l.
- Therefore, it can be a negative integer, zero,
or a positive integer.
15
3. Magnetic quantum
number (m)

16. Example: Fluorine (F)
- Fluorine has an { value of 1. Therefore, its
possible m, values are -1, 0, or 1.
- Using the Hund's rule, we can draw its five
valence electrons.
- Hund's rule states that every orbital in a
subshell must be singly occupied before it can
be doubly occupied.

18. Example: Fluorine (F)
- From the orbital diagram of fluorine, we can
see that it fills up to the 2p orbital, even if it is
only partially filled.
- If it has an electron on the 2p, orbital, it will
have an ml value of -1.

19. Example: Fluorine (F)
- If it has an electron on the 2p, orbital, it will
have an m, value of 0. Lastly, if it has an
electron on the 2p, orbital, it will have an m,
value of +1.
- Therefore, fluorine has an m, value of +1.

20. - This describes the spin of the electron. Its
values can be + 1/2 represented by an
upward arrow, or – 1/2 , represented by a
downward arrow.
- Electrons occupy each orbital singly first
before pairing up and filling the other
orbitals, as stated by the Hund's rule.
20
4. Electron spin quantum number
(m)

21. - The last inserted spin will be the quantum
spin of the atom.
- An orbital can accommodate only two
electrons with opposite spins in each
orbital.
21
4. Electron spin quantum number
(m)

22. Example: Fluorine (F)
- Since fluorine has only one electron in the 3p
orbital, it will have a "spin down" so its
electron spin quantum number is - ½.

23. A summary of the quantum numbers is given
below.

24. 24
Electron
Configuration

25. - The electron configuration of an atom is a
representation of its arrangement and
distribution.
- Electron configuration is used not only to
describe the orbitals of an atom in its
ground state, but also to represent cations
or anions by reflecting loss or gain of
electrons in their subsequent orbitals. 25
Electron Configuration

26. - Many of the physical and chemical
properties of elements can be explained by
their unique electron configuration.
- The valence electrons, or the electrons in
the outermost shell, determine the unique
chemistry of the element.
26
Electron Configuration

27. 27
Shells

28. - A simple way to show the arrangement of
electrons around an atom is to arrange the
electrons in energy levels or shells around the
nucleus of an atom.
- Electrons that are closest to the nucleus, or
those occupying the first energy level,
have the lowest energy.
- Electrons that are farther away from the
nucleus have higher energy. 28
Shells

29. 29
Orbitals

30. - Although electrons can be represented in
shells circling the nucleus of an atom, they
in fact, move along complicated paths.
- These paths are called orbitals or
subshells.
- There are four different orbital shapes: s,
p, d, and f.
30
Orbitals

31. 31

32. 32
Rules for
Assigning Orbitals

33. - The nucleus of an atom contains protons
and neutrons. Surrounding the nucleus
are the electrons.
- Electrons have the same negative charge;
they also have the same mass.
- However, each electron in an atom has a
different amount of energy.
33
Rules for Assigning Orbitals

34. - The electrons closest to the nucleus have
the lowest energy; the farthest have the
highest energy.
- Electrons in an atom fill the principal
energy levels in order of increasing energy.
- The order of levels filled is given below:
- 1s2
, 2s2
, 2p6
, 3s2
, 3p6
, 4s2
, 3d10
, 4p6
, 5s2
, 4d, 5p,
6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p 34
Rules for Assigning Orbitals

35. 35

36. 36
Hund's Rule

37. - When assigning electrons to orbitals, each
electron will first fill all the orbitals with the
same energy or degenerate orbitals before
pairing with another electron.
- Atoms at ground state tend to have as many
unpaired electrons as possible.
- When visualizing this process, think about how
electrons fill orbitals like most people do on
seats of an empty bus. 37
Hund's Rule

38. - A passenger would not sit next to a total
stranger if there are still empty rows
available.
- Only after all the seats are taken would a
passenger sit next to someone or pair up.
- This is why this rule is also often called, "the
empty bus seat rule."
38
Hund's Rule

39. 39
The Aufbau
Principle

40. - Aufbau is derived from the German word
aufbauen, which means "to build."
- Electron configurations are written in such
a way that orbitals are built up from atom
to atom.
- In writing the electron configuration of an
atom, orbitals are filled in order of
increasing atomic number 40
The Aufbau Principle

41. 41

42. 42
Writing Electron
Configurations

43. - An electron configuration is written
according to the following considerations:
• First, write the energy level (the period).
• Then, write the subshell to be filled.
• Then, write the superscript -- the number
of electrons in that subshell.
43
Writing Electron Configurations

44. - The total number of electrons is the atomic
number, Z.
- The rules above hold for writing the
electron configurations of all the elements
in the periodic table.
44
Writing Electron Configurations

45. - Three methods are used to write electron
configurations:
1. orbital diagram
2. spdf notation
3. noble gas notation
- Each method has its advantages and
disadvantages.
45
Writing Electron Configurations

46. 46
Orbital Diagram

47. - An orbital diagram is a representation of
electron configuration by which each of the
separate orbitals and the spins of the
electrons are shown.
- This is carried out by determining the
subshell (s, p, d, or f) and then drawing each
electron following the three considerations
in writing electron configurations. 47
Orbital Diagram

48. 48

49. 49
spdf Notation

50. - The spdf notation is perhaps the most
common way of writing electron
configuration ions.
- Unlike an orbital diagram, the distribution
of electrons in each orbital is not shown.
- The total number of electrons in each
energy level is indicated by a superscript.
50
spdf Notation

51. 51

52. 52
Noble Gas
Notation

53. - Noble gas notation is a simpler way of
writing electron configurations compared
to the spdf notation.
- Noble gases, which are also known as inert
gases, have the most stable electron
configurations.
- Because noble gases have their subshells
filled, they can be used as a simpler way of
writing electron configurations for
53
Noble Gas Notation

54. 54

55. 55
Magnetic
Properties

56. - Magnetism is a physical phenomenon
produced by the motion of electric charges.
- This can result in attraction or repulsion,
depending on the force between the atoms.
- It is also a property of a material to respond
to a particular magnetic field.
56
Magnetic Properties

57. 57
Paramagnetism

58. - Paramagnetism is the magnetic state of an
atom with one or more unpaired electrons.
- According to Hund's rule, each orbital must
be filled singly before it can be doubly
occupied.
- Thus, this results in unpaired electrons.
58
Paramagnetism

59. - The unpaired electrons can orient in either
direction causing the material's electrons to
align with a magnet.
- This allows paramagnetic atoms attracted
to magnetic fields.
59
Paramagnetism

60. 60

61. 61
Diamagnetism

62. - Diamagnetism, on the other hand, is the
magnetic state wherein an atom has no
unpaired electrons.
- As stated in Pauli's Exclusion Principle, no
two identical atoms can fill the same
quantum state/number at the same time.
62
Diamagnetism

63. - This causes opposite spins in a given orbital.
As a result, there is no magnetic field, and
the material is not attractive.
- In fact, diamagnetism causes a weak
repulsion.
63
Diamagnetism

64. 64
Practice Exercise

65. 65
Draw the orbital diagram of
the following using their
electronic configuration.

66. 66
1.Oxygen
2.Nitrogen
3.Flourine
4.Neon
5.Magnesium

67. 67

68. 68

69. 69

70. 70

71. 71

72. 72
Write the Electron
Configurations of the
following using Noble Gas
Notation:

73. 73
1.Boron
2.Aluminum
3.Gallium
4.Indium
5.Thalium

74. 74

75. 75

76. 76

77. 77

78. 78

79. 79
Generalization

80. Share one thing that
you’ve learned in
today’s discussion. 80
Generalization

81. Proceed to google
classroom and
answer the posted
activity. 81
ACTIVITY

82. Thanks!
82