Lesson 2, Classifying matter- Solid, Liquid, Gas and the others

1. composition
structure
Matter
properties

2. Matter Defined
2

3. • Matter is anything that has mass
and occupies space.
• Matter can be invisible.
– Air is matter, but it cannot be seen.
• Matter appears to be continuous and
unbroken.
– Matter is actually discontinuous. It is
made up of tiny particles call atoms.
3

4. An apparently empty test tube is submerged, mouth
downward in water. Only a small volume of water
rises into the tube, which is actually filled with
invisible matter–air.
4
3.1

5. Physical States of
Matter
5

6. 6

7. 7

8. Pure Substances and
Mixtures
8

9. Matter refers to all of the materials that
make up the universe.
9

10. Classification of matter: A pure substance is
always homogeneous in composition, whereas a
mixture always contains two or more substances
and may be either homogeneous or
heterogeneous.
3.2
10

11. Pure Substance
A particular kind of matter that has a fixed
composition and distinct properties.
Examples
ammonia, water, and oxygen.
11

12. Elements
12

13. An element is a fundamental or
elementary substance that cannot be
broken down into simpler substances by
any means
13

14. • All known substances on Earth and
probably the universe are formed by
combinations of more than 100
elements.
• Each element has a number.
– Beginning with hydrogen as 1 the
elements are numbered in order of
increasing complexity.
14

15. • Most substances can be decomposed
into two or more simpler substances.
– Water can be decomposed into hydrogen
and oxygen.
– Table salt can be decomposed into
sodium and chlorine.
• An element cannot be decomposed into
a simpler substance.
15

16. Distribution of
Elements
16

17. • Elements are not distributed equally
by nature.
– Oxygen is the most abundant element in
the human body (65%).
– Oxygen is the most abundant element in
the crust of the earth (49.2%)..
• In the universe the most abundant
element is hydrogen (91%) and the
second most abundant element is
helium (8.75%).
17

18. Distribution
of the
common
elements in
nature.
18
3.3

19. Names of the
Elements
19

20. Sources of Element Names
GreekColor
• Iodine: from the Greek iodes meaning
violet.
• Fluorine: from the Latin fluere meaning to
LatinProperty flow.
German- • Bismuth: from the German
Color
weisse mass which means white mass.
Location • Germanium: discovered in 1866 by a
German chemist.
Famous- • Einsteinium: named for Albert Einstein.
20
Scientists

21. Compounds
21

22. A compound is a distinct substance that
contains two or more elements combined
in a definite proportion by weight.
Compounds can be decomposed
chemically into simpler substances–that
is, into simpler compounds or elements.
22

23. There are two types of compounds:
molecular and ionic.
23

24. Molecules
24

25. A molecule is the smallest uncharged
individual unit of a compound formed by
the union of two or more atoms.
25

26. • A water molecule consists of two
hydrogen atoms and one oxygen atom.
• If it is subdivided the water molecule
will be destroyed and hydrogen and
oxygen will be formed.
26
3.5

27. Ionic Compounds
27

28. An ion is a positively or negatively
charged atom or group of atoms.
28

29. A cation is a positively charged ion.
29
3.5

30. An anion is a negatively charged ion.
30
3.5

31. Ionic compounds are held together by
attractive forces between positively and
negatively charged ions.
31

32. Sodium Chloride
The ultimate particles of sodium chloride
are positively charged sodium ions and
negatively charged chloride ions.
32
3.5

33. Compounds can be classified as molecular or
ionic. Ionic compounds are held together by
attractive forces between their positive and
negative charges. Molecular compounds are
held together by covalent bonds.
33
3.4

34. Mixture
Matter containing 2 or more substances that
are present in variable amounts. Mixtures
are variable in composition. They can be
homogeneous or heterogeneous.
34

35. Homogeneous Mixture (Solution)
A homogeneous mixture of 2 or more
substances. It has one phase.
Example
Sugar and water. Before the sugar and
water are mixed each is a separate phase.
After mixing the sugar is evenly dispersed
throughout the volume of the water.
35

36. Heterogeneous Mixture
A heterogeneous mixture consists of 2 or
more phases.
Example
Sugar and fine white sand. The amount of
sugar relative to sand can be varied. The
sugar and sand each retain their own
properties.
36

37. Heterogeneous Mixture
A heterogeneous mixture consists of 2 or more
phases.
Example
• Iron (II) sulfide (FeS) is 63.5% Fe and 36.5% S
by mass.
• Mixing iron and sulfur in these proportions does
not form iron (II) sulfide. Two phases are
present: a sulfur phase and an iron phase.
• If the mixture is heated strongly a chemical
reaction occurs and iron (II) sulfide is formed.
• FeS is a compound of iron and sulfur and has
37
none of the properties of iron and sulfur.

38. Mixture of iron
and sulfur
Compound of iron
and sulfur
Formula
Has no definite
formula: consists
of Fe and S.
FeS
Composition
Contains Fe and S
in any proportion
by mass.
63.5% Fe and
36.5% S by mass.
Separation
Fe and S can be
separated by
physical means.
Fe and S can be
separated only by
chemical change.
38

39. 39

40. Symbols of the
Elements
40

41. • A symbol stands for
– the element itself
– one atom of the element
– a particular quantity of the element
41

42. Rules governing symbols of the elements
are:
1. Symbols have either one or
two letters.
2. If one letter is used it is
capitalized.
C hydrogen
H carbon
3. If two letters are used, only Ne barium
Ba neon
the first is capitalized.
42

43. These symbols have carried over from the earlier names of the
A number of symbols Latin). to the sameconnection with the element.
elements (usually start with have no letter as the element.
Most symbols appear
43

44. 44

45. Metals, Nonmetals
and Metalloids
45

46. Metals
46

47. Most elements
are metals
47

48. • Metals are solid at room temperature.
– Mercury is an exception. At room temperature it
is a liquid.
• Metals are good conductors of heat and electricity.
Most elements
are metals
physical
properties
of or
• Metals are malleable (they can be rolledmetals
hammered into sheets).
• Metals have high luster (they are shiny).
48

49. • Metals are ductile (they can be drawn into wires).
• Most metals have a high melting point.
Most elements
are metals
• Metals have high densities
49

50. Examples of Metals
lead
gold
iron
50

51. Chemical Properties of Metals
• Metals have little tendency to combine with
each other to form compounds.
• Many metals readily combine with nonmetals
to form ionic compounds.
– They can combine with sulfur.
oxygen.
chlorine.
– In nature, minerals are formed by combinations
of the more reactive metals combined with other
elements.
51

52. Chemical Properties of Metals
– A few of the less reactive metals such as copper,
silver and gold are found in the free state.
– Metals can mix with each other to form alloys.
 Brass is a mixture of copper and zinc.
 Bronze is a mixture of copper and tin.
 Steel is a mixture of carbon and iron.
52

53. Nonmetals
53

54. Physical Properties of Nonmetals
• Lack luster (they are dull)
• Have relatively low melting points
• Have low densities.
• Poor conductors of heat and electricity
• At room temperature carbon, phosphorous,
sulfur, selenium, and iodine are solids.
54

55. Physical State at Room Temperature
phosphorous
carbon
Solid
selenium
sulfur
iodine
55

56. Physical State at Room Temperature
liquid
bromine
56

57. Physical State at Room Temperature
nitrogen,
oxygen
gas
fluorine,
chlorine
helium, neon, argon, krypton, xenon, radon
57

58. Metalloids
58

59. Metalloids have properties that
are intermediate between metals
and nonmetals
59

60. The Metalloids
1. boron
2. silicon
3. germanium
4. arsenic
5. antimony
6. tellurium
7. polonium
60

61. Nonmetals found to to left of the the metalloids.
Metals are are found thethe right of metalloids
61

62. In 1869 Dimitri Mendeleev of Russia and
Lothar Meyer of Germany independently
published periodic arrangements of the
elements based on increasing atomic
masses.
Mendeleev’s arrangement is the precursor
to the modern periodic table.
62

63. Period numbers correspond
Horizontal rows are
to the highest occupied
called periods.
energy level.
63
10.14

64. Elements in the A groups
Elements in the B groups
with similar
Groups are numbered
are designated organized
properties are transition
are designated
with Roman numerals.
in groups or families.
representative elements.
elements.
64
10.14

65. Period number corresponds with the
highest energy level occupied by
electrons in that period.
65
10.17

66. The group numbersfamily have the same
The elements of a for the representative
outermost electron configurationnumber that
elements are equal to the total except of
outermost electrons in the atoms of the group.
the electrons are in different energy levels.
66
10.17

67. Noble Gas
Halogen
Group
Alkali Metal
Alkali Earth Metal
Period

68. Periodic Trends in
Atomic Properties
68

69. Characteristic properties and trends of the
elements are the basis of the periodic
table’s design.
69

70. These trends allow us to use the periodic
table to accurately predict properties and
reactions of a wide variety of substances.
70

71. Atomic Radius
71

72. Radii of atoms
increase down a
group.
For each step down a group, electrons enter
the next higher energy level.
72
11.2

73. Radii of atoms tend to decrease
from left to right across a period.
This
Each time an
For increase in
positive is added
electron nuclear
representative a
charge pulls all
proton is within
elements a added to
electrons closer
the same period to
nucleus.
the energy level
nucleus.
remains constant
as electrons are
added.
73
11.2

74. Radii of atoms tend to decrease
from left to right across a period.
Each
This time an
For increase in
electron is added
positive nuclear
representative a
proton pulls all
charge is within
elements added to
electrons closer
the same period to
nucleus.
the energy level
nucleus.
remains constant
as electrons are
added.
74
11.2

75. Ionization Energy
75

76. The ionization energy of an atom is the
energy required to remove an electron from
an atom.
Na + ionization energy → Na+ + e-
76

77. As each succeeding electron is removed from
an atom ever higher energies are required.
77

78. The Atom
Atom is the basic unit of an element,
made up of even smaller particles
called subatomic particles.
There are three fundamental
components (subatomic particles) that
are important in chemistry: Electron,
Proton and Neutron.
The protons and neutrons of an atom
are packed in an extremely small
nucleus.
Electrons are shown as ‘clouds’ around
the nucleus.

79. The Structure of the Atom
The Structure of the Atom
Electron (cloud)
Nucleus
Figure above shows the location of the protons,
Neutrons and electrons in an atom

80. Subatomic Particles
mass p = mass n = 1840 x mass e-

81. Atomic Numbers of
the Elements
81

82. Atomic Number, Mass Number & Isotopes
Atomic number (Z)
= number of protons in nucleus
Mass number (A)
= number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different numbers of
neutrons in their nuclei
Mass Number
Atomic Number
A
ZX
Element Symbol

83. • The atomic number of an element is
equal to the number of protons in the
nucleus of that element.
• The atomic number of an atom
determines which element the atom is.
83

84. Every atom with an atomic
number of 1 is a hydrogen atom.
Every hydrogen atom contains 1
proton in its nucleus.
84

85. Every atom with an atomic
number of 6 is a carbon atom.
Every carbon atom contains 6
protons in its nucleus.
85

86. atomic
number
Every atom with an
atomic number of 1
is a hydrogen atom.
1 proton in the
nucleus
H
1
86

87. atomic
number
Every atom with an
atomic number of 6
is a carbon atom.
6 protons in the
nucleus
C
6
87

88. atomic
number
U
92
Every atom with an
atomic number of
92 is a uranium
atom.
92 protons
in the
nucleus
88

89. Isotopes of the
Elements
89

90. • Atoms of the same element can have
different masses.
• They always have the same number of
protons, but they can have different
numbers of neutrons in their nuclei.
• The difference in the number of neutrons
accounts for the difference in mass.
• These are isotopes of the same element.
90

91. Isotopic Notation
91

92. Isotopic Notation
6 protons + 6 neutrons
12
C
6
6 protons
92

93. Isotopic Notation
6 protons + 8 neutrons
14
C
6
6 protons
93

94. Isotopic Notation
8 protons + 8 neutrons
16
O
8
8 protons
94

95. Isotopic Notation
8 protons + 9 neutrons
17
O
8
8 protons
95

96. Isotopic Notation
8 protons + 10 neutrons
18
O
8
8 protons
96

97. Hydrogen has three isotopes
1 proton
1 proton
1 proton
0 neutrons
1 neutron
2 neutrons
97

98. Relationship Between Mass
Number and Atomic Number
98

99. The mass number minus the atomic
number equals the number of neutrons in
the nucleus.
mass
number
atomic
number
atomic
mass number number
109
47
=
=
number of
neutrons
62
99

100. Energy Levels
of Electrons
100

101. As n increases, the
energy of the electron
increases.
The first four
principal energy
levels of an atom.
Each level is
assigned a principal
quantum number n.
101
10.7

102. 10.7, 10.8
Each principal energy level
is subdivided into sublevels.
102

103. Within sublevels the electrons are found in
orbitals.
An s orbital is spherical in
shape.
The
spherical
surface
encloses a space where
there is a 90% probability
that the electron may be
found.
103
10.10

104. An atomic orbital can hold a maximum of two
electrons.
An electron can spin in one
of two possible directions
represented by ↑ or ↓.
The two electrons that
occupy an atomic orbital
must have opposite spins.
This is known as the Pauli
Exclusion Principal.
104
10.10

105. A p sublevel is made up of three orbitals.
Each p orbital has two lobes.
Each p orbital can hold a maximum of two
electrons.
A p sublevel can hold a maximum of 6
electrons.
10.10
105

106. pz
The three p orbitals share
a common center.
py
px
The three p orbitals point
in different directions.
106
10.10

107. A d sublevel is made up of five orbitals.
The five d orbitals all point in different directions.
Each d orbital can hold a maximum of two
electrons.
A d sublevel can hold a maximum of 10
electrons.
10.11
107

108. 10.8 10.10 10.11
Number of Orbitals in a Sublevel
108

109. Distribution of Subshells by
Principal Energy Level
n=1
1s
n=2
2s
2p 2p 2p
n = 3 3s
3p 3p 3p
3d 3d 3d 3d 3d
n = 4 4s
4p 4p 4p
4d 4d 4d 4d 4d
4f 4f 4f 4f 4f 4f 4f
109

110. Nuclear makeup and electronic structure of
each principal energy level of an atom.
number of protons and electrons
number of
neutrons in thein each sublevel
nucleus
110
10.13

111. Electron Configuration
Number of
electrons in
sublevel orbitals
Arrangement of
electrons within their
respective sublevels.
2p
6
Principal
Type of orbital
energy level
111

112. The electron configuration of any of the
noble gas elements can be represented by
the symbol of the element enclosed in
square brackets.
B
1s22s22p1
[He]2s22p1
Na
1s22s22p63s1
[Ne]3s1
Cl
1s22s22p63s23p5
[Ne]3s23p5
112

113. The electron configuration of argon is
Ar 1s22s22p63s23p6
The elements after argon are potassium
and calcium. Instead of entering a 3d
orbital, the valence electrons of these
elements enter the 4s orbital.
K
1s22s22p63s23p64s1
[Ar]4s1
Ca
1s22s22p63s23p6 4s2
[Ar]4s2
113

114. d orbital numbers are 1 less
than d orbital filling
the period number
10.16
Arrangement of electrons
according to sublevel being filled.
114

115. f orbital numbers are 2 less
than fthe period number
orbital filling
10.16
Arrangement of electrons
according to sublevel being filled.
115