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1. 1
GENERAL CHEMISTRY
AND
INORGANIC CHEMISTRY

Dr. BÙI THỊ BỬU HUÊ
College of Natural Science

2. 2
Chapter 1. ATOMIC STRUCTURE AND THE
PERIODIC TABLE
Chapter 2. CHEMICAL BONDS AND MOLECULAR
STRUCTURE
Chapter 3. CHEMICAL THERMODYNAMICS
Chapter 4. CHEMICAL KINETICS
Chapter 5. CHEMICAL EQUILIBRIUM
Chapter 6. SOLUTIONS
Chapter 7. ACIDS AND BASES
Chapter 8. CHEMISTRY OF METALS
Chapter 9. CHEMISTRY OF NONMETALS
Chapter 10. TRANSITION METALS AND COMPLEXES

3. References
1. Brady and Holum, 1996, Chemistry: the Study
of Matter and its Changes,
2
th
Ed., John Wiley & Sons. Inc. New York.
2. Umland, Jean B., 1993, General Chemistry,
West publishing company.
3. Zumdahl, Steven S., 1995, Chemical Principal,
2
th
Ed. DC. Health & company. Toronto.
4. http://www.chemistry.msu.edu/Courses/
5. http://antoine.frostburg.edu
6. http://chemed.chem.purdue.edu
7. http://www.chem1.com/chemed/genchem.html
8. http://www.cbu.edu/~mcondren/lectures.htm
9. http://ull.chemistry.uakron.edu/GenChem/index.html

4. Chapter 1. ATOMIC STRUCTURE AND
THE PERIODIC TABLE
➢ Understand atomic structure of an atom
including its mass number, isotopes and
orbitals.
➢ Know how to account for the structure
of the periodic table of the elements based
on the modern theory of atomic structure.
➢ Understand general trends of several
important atomic properties.

5. FUNDAMENTAL PARTICLES
An atom is composed of three types of
subatomic particles: the proton, neutron,
and electron
ParticleMass (g)Charge
Proton 1.6727 x 10
-24
+1
Neutron1.6750 x 10
-24
0
Electron9.110 x 10
-28
-1

6. 6
Atomic Structure
Atoms consist of very small, very dense positively
charged nuclei surrounded by clouds of electrons at
relatively great distances from the nuclei.

7. 7
Nuclide Symbol
Mass number = number of protons + number of neutrons
= atomic number + neutron number

8. 8
Isotopes are atoms of the same element
with different masses; they are atoms
containing the same number of protons but
different numbers of neutrons
ISOTOPES

9. 9
The three isotopes of Hydrogen
Particle Mass (g) Charge
Proton 1.6727 x 10
-24
+1
Neutron 1.6750 x 10
-24
0
Electron 9.110 x 10
-28
-1
1 amu = 1.660 x 10
-24
g

10. 10
THE ATOMIC WEIGHT SCALE AND
ATOMIC WEIGHTS
➢ The atomic weight scale is based on
the mass of the carbon-12 isotope
➢ One amu is exactly 1/12 of the mass
of a carbon-12 atom:
1 g = 6.022 x 10
23
amu
or 1 amu = 1.660 x 10
-24
g

11. 11
Atomic weight =0.7899 (23.98504 amu) +
0.1000 (24.98584 amu) +
0.1101 (25.98259 amu)
=18.946 amu +2.4986 amu +2.8607 amu = 24.30 amu
THE ATOMIC WEIGHT SCALE AND
ATOMIC WEIGHTS

12. 12
ELECTRONIC STRUCTURES OF ATOMS
✓ Why do different elements have such different
chemical and physical properties?
✓ Why does chemical bonding occur at all?
✓ Why does each element form compounds
with characteristic formulas?
✓ How can atoms of different elements give
off or absorb light only of characteristic
colors.

13. 13

14. 14
Electromagnetic Radiation
= c
Where:
 frequency
 wavelength
c: speed of light
c = 2.99 x 10
8
m/s

15. 15
Electromagnetic Radiation

16. 16
Photons
The quantum of electromagnetic energy,
generally regarded as a discrete particle having
zero mass, no electric charge, and an
indefinitely long lifetime.
E = hν = hc/λ
h = Planck's constant = 6.626 × 10
−34
J.s

17. 17

18. 18

19. 19
Electromagnetic Spectrum

20. 20
Dispersion of White Light

21. 21
EMISSION & ABSORPTION SPECTRA

22. 22
ATOMIC SPECTRA

23. 23
In 1913, Niels Bohr (1885–1962):
▪The electronic energy is quantized:
only certain values of electronic
energy are possible.
▪ The electrons absorb or emit energy
in discrete amounts as they move from
one orbit to another.
Bohr Model

24. 24
Bohr Model

25. 25
Bohr Model for Hydrogen Atom
▪ Allowed orbits:
mvr = nh/2
r = n
2
a
0
a
0= 5.292 x 10
−11
m = 0.5292 Å
n = quantum number
= 1, 2, 3, 4, 5, 6, etc
▪ Potential energy:

26. 26
Ground State
The state of least possible energy in a
physical system, as of elementary
particles. Also called ground level.

27. 27
Excited State
Being at an energy level higher than
the ground state.

28. 28
Electron Transition in a Hydrogen Atom
Lyman series → ultraviolet
n > 1 → n = 1
Balmer series → visible light
n > 2 → n = 2
Paschen series → infrared
n > 3 → n = 3

29. 29

30. 30

31. 31
Quantum Mechanics
Theory of the structure and behavior of
atoms and molecules.

32. 32

33. 33

34. 34

35. 35

36. 36

37. 37
The Schrödinger Equation
▪ Has been solved exactly only for one-
electron species such as the hydrogen atom
and the ions He
+
and Li
2+
.
▪ Simplifying assumptions are necessary to
solve the equation for more complex atoms
and molecules.

38. 38
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).
Atoms and molecules possess various forms of
energy. Let us focus our attention on their electronic
energies.
Basic Ideas of Quantum Mechanics

39. 39
2. When atoms or molecules emit or absorb
radiation (light), they change their energies.
The energy change in the atom or molecule is
related to the frequency or wavelength of the light
emitted or absorbed by the equations:
ΔE = hν or ΔE = hc/λ
The energy lost (or gained) by an atom as it goes
from higher to lower (or lower to higher) energy
states is equal to the energy of the photon emitted
(or absorbed) during the transition.

40. 40
3. The allowed energy states of atoms and molecules
can be described by sets of numbers called quantum
numbers: n, l, m

41. 41

42. 42

43. 43

44. 44
Atomic Orbitals
▪An atomic orbital is a region of
space around the nucleus in which
the probability of finding an
electron is high.
▪Determined by a set of quantum
numbers: n, l, m.
▪4 types: s, p, d, f.

45. 45
Atomic Orbitals, s-type

46. 46
Atomic Orbitals, p-type

47. 47
Atomic Orbitals, d-type

48. 48

49. 49

50. 50
Electronic Configurations
•The shorthand representation of the occupancy
of the energy levels (shells and subshells) of an
atom by electrons.

51. 51

52. 52

53. 53

54. 54

55. 55

56. 56

57. 57

58. 58

59. 59
Hund's Rules

60. 60
Electronic Configuration
H atom (1 electron): 1s
1
He atom (2 electrons): 1s
2
Li atom (3 electrons): 1s
2
, 2s
1
Cl atom (17 electrons): 1s
2
, 2s
2
, 2p
6
, 3s
2
, 3p
5

61. 61
Electronic Configuration
As atom (33 electons):
1s
2
, 2s
2
, 2p
6
, 3s
2
, 3p
6
, 4s
2
, 3d
10
, 4p
3
or
[Ar] 4s
2
, 3d
10
, 4p
3

62. 62

63. 63

64. 64
Electronic Configuration
Negative ions:
add electron(s), 1 electron for each
negative charge
S
-2
ion: (16 + 2) electrons:
1s
2
, 2s
2
, 2p
6
, 3s
2
, 3p
6

65. 65
Electronic Configuration
Positive ions
remove electron(s), 1 electron for each
positive charge
Mg
+2
ion: (12-2) electrons:
1s
2
, 2s
2
, 2p
6

66. 66
How many valence electrons are in Cl:
[Ne]3s
2
3p
5
?
2, 5, 7

67. 67
For Cl to achieve a noble gas configuration,
it is more likely that:
electrons would be added
electrons would be removed

68. 68

69. 69
Regions by Electron Type

70. 70

71. 71

72. 72
Trends in the Periodic Table
• Atomic radius
• Ionic radius
• Ionization energy
• Electron affinity

73. 73
Atomic Radius
decrease left to right across a period
Z
eff = Z - S
where:
Z
eff = effective nuclear charge
Z = nuclear charge, atomic number
S = shielding constant

74. 74
▪Increase top to bottom down a group
▪Increases from upper right corner to
the lower left corner
Atomic Radius

75. 75
Atomic Radius

76. 76
Atomic Radius vs. Atomic Number

77. 77
Ionic Radii

78. 78
•Same trends as for atomic radius
Ionic Radii

79. 79
Comparison of Atomic and Ionic Radii
▪ Positive ions smaller than atom
▪ Negative ions larger than atom

80. 80
Isoelectronic Series
•series of negative ions, noble gas
atom, and positive ions with the
same electronic confiuration
•size decreases as “positive charge”
of the nucleus increases
Ionic Radii

81. 81
Ionization Energy
•energy necessary to remove an electron to
form a positive ion
•low value for metals, electrons easily
removed
•high value for non-metals, electrons difficult
to remove
•increases from lower left corner of periodic
table to the upper right corner

82. 82
first ionization energy
Energy to remove first electron from an
atom.
second ionization energy
Energy to remove second electron from a
+1 ion, etc.
Ionization Energy

83. 83
Ionization Energy vs. Atomic Number

84. 84
Electron Affinity
•Energy released when an electron is
added to an atom
•Same trends as ionization energy,
increases from lower left corner to the
upper right corner
•Metals have low “EA”
•Nonmetals have high “EA”

85. 85
Magnetism
•Result of the spin of electrons
•Diamagnetism: no unpaired electrons
•Paramagnetism: one or more unpaired
electrons

86. 86

87. 87