Chemistry-ADGE108-cl2024-2025-Unit-5.ppt

Notes

1. 1
Chemical Bonds:
The Formation of
Compounds From Atoms

2. 2
Periodic Trends in Periodic Trends in
Atomic PropertiesAtomic Properties

3. 3
Characteristic properties and trends of the
elements are the basis of the periodic
table’s design.

4. 4
These trends allow us to use the periodic
table to accurately predict properties and
reactions of a wide variety of substances.

5. 5
Metals and NonmetalsMetals and Nonmetals

6. 6
Chemical Properties
of Metals
•metals tend to lose
electrons and form
positive ions.
•nonmetals tend to
gain electrons and
form negative ions.
Chemical Properties of
Nonmetals
When metals react with nonmetals electrons
are usually transferred from the metal to the
nonmetal.

7. 7
Physical Properties
of Metals
•lustrous
•malleable
•good conductors of
heat
•good conductors of
electricity
•nonlustrous
•brittle
•poor conductors of
heat
•poor conductors of
electricity
Physical Properties
of Nonmetals

8. 8
Metalloids have properties that
are intermediate between metals
and nonmetals

9. 9
The Metalloids
1.boron
2.silicon
3.germanium
4.arsenic
5.antimony
6.tellurium
7.polonium

10. 10
Metals are found to the left of the metalloidsNonmetals are found to the right of the metalloids.
11.1

11. 11
Atomic RadiusAtomic Radius

12. 12
Radii of atoms
increase down a
group.
For each step down a group, electrons enter
the next higher energy level.
11.2

13. 13
Radii of atoms tend to decrease
from left to right across a period.
For
representative
elements within
the same period
the energy level
remains constant
as electrons are
added.
Each time an
electron is added a
proton is added to
the nucleus.
This increase in
positive nuclear
charge pulls all
electrons closer to
the nucleus.
11.2

14. 14
Radii of atoms tend to decrease
from left to right across a period.
For
representative
elements within
the same period
the energy level
remains constant
as electrons are
added.
This increase in
positive nuclear
charge pulls all
electrons closer to
the nucleus.
11.2
Each time an
electron is added a
proton is added to
the nucleus.

15. 15
Ionization EnergyIonization Energy

16. 16
The ionization energy of an atom is the
energy required to remove an electron from
an atom.
Na + ionization energy → Na
+
+ e
-

17. 17
•The first ionization energy is the amount
of energy required to remove the first
electron from an atom.
He + first ionization energy → He
+
+ e
-

He + 2,372 kJ/mol → He
+
+ e
-

•The second ionization energy is the
amount of energy required to remove the
second electron from an atom.
He+ + 5,247 kJ/mol → He
++
+ e-
He + second ionization energy → He
+
+ e
-

18. 18
As each succeeding electron is removed from
an atom ever higher energies are required.

19. 19
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
11.3
Ionization energies gradually increase from left to
right across a period.
IA
IIA
IIIA
IVA
VA
VIA
VIIA
Noble
Gases
1
2
3
4

20. 20
Periodic relationship of the first ionization energy for
representative elements in the first four periods.
11.3
Ionization energies of Group A elements decrease
from top to bottom in a group.
IA
IIA
IIIA
IVA
VA
VIA
VIIA
Noble
Gases
D
is
t
a
n
c
e

o
f

O
u
t
e
r

S
h
e
ll
E
le
c
t
r
o
n
s

F
r
o
m

N
u
c
le
u
s

nonmetals
metals
nonmetals have higher ionization
potentials than metals

21. 21
Lewis StructuresLewis Structures
of Atomsof Atoms

22. 22
Metals form cations and nonmetals form
anions to attain a stable valence electron
structure.

23. 23
This stable structure often consists of two s and
six p electrons.
These rearrangements occur by losing, gaining,
or sharing electrons.

24. 24
• Na with the electron structure 1s
2
2s
2
2p
6
3s
1
has 1 valence electron.
The Lewis structure of an atom is a
representation that shows the valence
electrons for that atom.
• Fluorine with the electron structure 1s
2
2s
2
2p
5
has 7 valence electrons

25. 26
The Lewis structure of an atom uses dots to
show the valence electrons of atoms.
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
B Symbol of
the element
2s
2
2p
1

26. 27
The number of dots equals the number of s
and p electrons in the atom’s outermost shell.
S
2s
2
2p
4
The Lewis structure of an atom uses dots to
show the valence electrons of atoms.

27. 28
11.4
Lewis Structures of the first 20 elements.

28. 29
The Ionic Bond: Transfer ofThe Ionic Bond: Transfer of
Electrons From One AtomElectrons From One Atom
to Anotherto Another

29. 30
The chemistry of many elements,
especially the representative ones, is to
attain the same outer electron structure as
one of the noble gases.

30. 31
With the exception of helium, this structure consists
of eight electrons in the outermost energy level.

31. 32
After sodium loses its 3s electron it has attained the
same electronic structure as neon.

32. 33
After chlorine gains a 3p electron it has attained the
same electronic structure as argon.

33. 34
Formation of NaCl

34. 35
The 3s electron of sodium transfers to the half-filled 3p
orbital of chlorine.
Lewis representation of sodium chloride formation.
A sodium ion (Na+) and a chloride ion (Cl
-
) are formed.
The force holding Na
+
and Cl
-
together is an ionic bond.

35. 36
Formation of MgCl
2

36. 37
Two 3s electrons of magnesium transfer to the half-filled
3p orbitals of two chlorine atoms.
A magnesium ion (Mg
2+
) and two chloride ions (Cl
-
) are
formed.
The forces holding Mg
2+
and two Cl
-
together are ionic
bonds.

37. 38
NaCl is made up of cubic crystals.In the crystal each sodium ion is surrounded by six
chloride ions.

38. 39
In the crystal each chloride ion is surrounded by six
sodium ions.
11.5

39. 40
The ratio of Na
+
to Cl
-
is 1:1
11.5

40. 41
Relative Size of
Sodium Ion to Chloride Ion

41. 42
Sodium ion

is smaller than a sodium atom
because:
(1) the sodium atom has lost its outermost
electron.
(2) The 10 remaining electrons are now
attracted by 11 protons and are drawn closer
to the nucleus.
11.6

42. 43
Chloride ion

is larger than a chloride atom
because:
(1) the chlorine atom has gained an electron
and now has 18 electrons and 17 protons.
(2) The nuclear attraction on each electron is
thereby decreased, allowing the chlorine to expand.
11.6

43. 44
•Metals usually have one, two or three
electrons in their outer shells.
•When a metal reacts it:
–usually loses one two or three electrons
–attains the electron structure of a noble
gas
–becomes a positive ion.
•The positive ion formed by the loss of
electrons is much smaller than the
metal atom.

44. 45
•Nonmetals usually have one, two or three
electrons in their outer shells.
•When a nonmetal reacts it:
–usually gains one two or three electrons
–attains the electron structure of a noble
gas
–becomes a negative ion.
•The negative ion formed by the gain of
electrons is much larger than the
nonmetal atom.

45. 46

46. 47
Predicting Formulas ofPredicting Formulas of
Ionic CompoundsIonic Compounds

47. 48
In almost all stable chemical compounds of
representative elements, each atom attains a
noble gas electron configuration.

48. 49
•Metals will lose electrons to attain a
noble gas configuration.
•Nonmetals will gain electrons to attain a
noble gas configuration.
Barium and Sulfur Combine.
–sulfur gains two electrons from barium and
attains an argon configuration.
–barium loses two electrons to sulfur and
attains a xenon configuration.
S [Ne]3s
2
3p
4
Ba [Xe]6s
2
Ba → [Xe] + 2e
-
S + 2e
-
→ [Ar]

Ba + S → BaS

49. 50
Because of similar electron structures,
the elements of a family generally form
compounds with the same atomic ratios.

50. 51

51. 52
10.17
The elements of a family have the same
outermost electron configuration except that
the electrons are in different energy levels.

52. 53
•The atomic ratio of the alkali metal
sodium to chlorine is 1:1 in NaCl.
•The atomic ratios of the other alkali
metal chlorides can be predicted to also
be 1:1.
•LiCl, KCl, CsCl, FrCl

53. 54
•The atomic ratio of hydrogen to nitrogen
is 3:1 in ammonia (NH
3). Nitrogen is the
first member of group VA.
•The atomic ratio of hydrogen when
combined with other group VA
elements can be predicted to also be
3:1.
•PH
3, AsH
3, SbH
3, BiH
3

54. 55
The Covalent Bond:The Covalent Bond:
Sharing ElectronsSharing Electrons

55. 56
A covalent bond consists of a pair of
electrons shared between two atoms.
In the millions of chemical compounds
that exist, the covalent bond is the
predominant chemical bond.

56. 57
Substances which covalently bond exist
as molecules.
Carbon dioxide bonds covalently.
It exists as individually bonded
covalent molecules containing
one carbon and two oxygen
atoms.
11.7

57. 58
The term molecule is not used when
referring to ionic substances.
Sodium chloride bonds ionically.
It consists of a large aggregate of
positive and negative ions. No
molecules of NaCl exist.
11.7

58. 59
Covalent bonding in the hydrogen molecule
Two 1s orbitals from each of
two hydrogen atoms overlap.
Each 1s orbital contains 1
electron.
The orbital of the
electrons includes
both hydrogen
nuclei.
The most likely
region to find the two
electrons is between
the two nuclei.
The two nuclei are
shielded from each
other by the electron
pair. This allows the
two nuclei to draw
close together.
Two 1s orbitals from each of
two hydrogen atoms overlap.
11.8

59. 60
11.9
Covalent bonding in the chlorine molecule
Each unpaired 3p orbital on
each chlorine atom contains
1 electron.
Two 3p orbitals from each of
two chlorine atoms overlap.
The orbital of the
electrons includes
both chlorine
nuclei.
The most likely
region to find the
two electrons is
between the two
nuclei.
The two nuclei are
shielded from each
other by the
electron pair. This
allows the two
nuclei to draw close
together.
Two 3p orbitals from each of
two chlorine atoms overlap.
Each chlorine now has 8
electrons in its outermost
energy level.

60. 61
hydrogen chlorine iodine nitrogen
Covalent bonding with equal sharing of
electrons occurs in diatomic molecules
formed from one element.
A dash may replace a pair of dots.

61. 62
ElectronegativityElectronegativity

62. 63
electronegativity The relative attraction
that an atom has for a pair of shared
electrons in a covalent bond.

63. 64
•If the two atoms that constitute a
covalent bond are identical then there
is equal sharing of electrons.
•This is called nonpolar covalent
bonding.
•Ionic bonding and nonpolar covalent
bonding represent two extremes.

64. 65
•If the two atoms that constitute a
covalent bond are not identical then
there is unequal sharing of electrons.
•This is called polar covalent bonding.
•One atom assumes a partial positive
charge and the other atom assumes a
partial negative charge.
–This charge difference is a result of the
unequal attractions the atoms have for
their shared electron pair.

65. 66
:
HCl
+ -
Shared electron pair.
:
The shared electron pair
is closer to chlorine than
to hydrogen.
Partial positive charge
on hydrogen.
Partial negative charge
on chlorine.
Chlorine has a greater attraction for the
shared electron pair than hydrogen.
Polar Covalent Bonding in HCl
The attractive force that an atom of an element has for
shared electrons in a molecule or a polyatomic ion is
known as its electronegativity.

66. 67
A scale of relative electronegativities
was developed by Linus Pauling.

67. 68
Electronegativity decreases down a group for
representative elements.
Electronegativity generally increases left to right
across a period.

68. 69
The electronegativities of the metals are low.
The electronegativities of the nonmetals are high.
11.1

69. 70
The polarity of a bond is determined by the
difference in electronegativity values of the
atoms forming the bond.

70. 71
•If the electronegativity difference
between two bonded atoms is greater
than 1.7-1.9 the bond will be more ionic
than covalent.
•If the electronegativity difference is
greater than 2, the bond is strongly
ionic.
•If the electronegativity difference is
less than 1.5, the bond is strongly
covalent.

71. 72
HH
Hydrogen Molecule
If the electronegativities are the same, the bond is
nonpolar covalent and the electrons are shared
equally.
The molecule is
nonpolar covalent.
Electronegativity
2.1
Electronegativity
2.1
11.10

72. 73
If the electronegativities are the same, the bond is
nonpolar covalent and the electrons are shared
equally.
Cl Cl
Chlorine Molecule
Electronegativity
3.0
Electronegativity
3.0
The molecule is
nonpolar covalent.
Electronegativity
Difference = 0.0
11.10

73. 74
If the electronegativities are not the same, the bond
is polar covalent and the electrons are shared
unequally.
H Cl
Hydrogen Chloride Molecule
Electronegativity
2.1
Electronegativity
3.0
The molecule is
polar covalent.
+ -
Electronegativity
Difference = 0.9
11.10

74. 75Sodium Chloride
Na
+
Cl
-
If the electronegativities are very different, the bond
is ionic and the electrons are transferred to the
more electronegative atom.
Electronegativity
0.9
Electronegativity
3.0
The bond is ionic.No molecule exists.
Electronegativity
Difference = 2.1
11.10

75. 76
A dipole is a molecule that is
electrically asymmetrical, causing it to
be oppositely charged at two points.
A dipole can be written as+ -

76. 77
An arrow can be used to indicate a dipole.
The arrow points to the negative end of the
dipole.
HClHBrH
O
H
Molecules of HCl, HBr and H
2
O are polar .

77. 78
A molecule containing different kinds of
atoms may or may not be polar depending
on its shape.
The carbon dioxide molecule is nonpolar
because its carbon-oxygen dipoles cancel
each other by acting in opposite directions.

78. 79
11.11
Relating Bond Type to
Electronegativity Difference.

79. 80
Lewis Structures ofLewis Structures of
CompoundsCompounds

80. 81
In writing Lewis structures, the most
important consideration for forming a
stable compound is that the atoms attain
a noble gas configuration.

81. 82
•The most difficult part of writing
Lewis structures is determining the
arrangement of the atoms in a molecule
or an ion.
•In simple molecules with more than
two atoms, one atom will be the central
atom surrounded by the other atoms.

82. 83
Cl
2
O has two possible arrangements.
Cl-Cl-O
The two chlorines can be bonded to each other.
Cl-O-Cl
The two chlorines can be bonded to oxygen.
Usually the single atom will be the central atom.

83. 84
Procedures for Writing
Lewis Structures

84. 85
Atom Group
Valence
Electrons
Cl VIIA 7
H IA 1
C IVA 4
N VA 5
S VIA 6
P VA 5
I VIIA 7
Valence Electrons of Group A Elements

85. 86
Step 1 Determine the number of electrons
needed to satisfy the octet rule for each
atom.
. Let us consider CCl
4
. Since this
compound has 4 chlorine atoms and 1
carbon atom, it has a total of 5 atoms that
need 8 valence electrons. Thus the total
number of electrons needed is 40.

86. 87
Step 2 Obtain the total number of valence
electrons to be used in the structure by
adding the number of valence electrons in
all the atoms in the molecule or ion.
If there is a charge, subtract an electron if it is a
positive charge, and add an electron if it is a
negative charge. For CCl
4 , we have:
Cl = 7 x 4 = 28
C = 4 x 1 = 4
Therefore, it has a total of 32 available valence electrons.

87. 88
Step 3 Determine the difference between the required electrons
and the number of valence electrons. The answer will be the
number of electrons involved in bond formation. Half of this
number will be the number of bonds used in the compound.
For CCl
4, the difference is 8 (40-32 = 8). This
means that CCl
4
has 4 bonds.

88. 89
Step 4 Draw the molecules with the
number of bonds you have calculated.
Below are some guidelines:
•If one atom is different from the rest, it is possible that it is the
central atom. Thus, in our example, chlorine atoms surround the
carbon atom.
•The carbon group (Group IVA) usually has four bonds. The
nitrogen group (Group VA) has three bonds, and the oxygen group
(Group VIA) has two bonds. Moreover, the halogens have one
bond.
•When oxygen and hydrogen are represented in a molecule, they
usually have the arrangement H-O-X, where X is another type of
element
•Remember that there may be double bonds and even triple bonds.

89. 90
Step 5 Fill in the dots that will represent
electron pairs. The bonds signify shared
electrons.
In CCl
4, there are 8 shared electrons. Therefore,
there are 6 electron pairs for each Cl atom.
Carbons does not have any dots around it since
the octet rule was satisfied already by the bonds.

90. 91
C
Cl
Cl
Cl
Cl::
:
:: :
::
::
:
:
:
:
:
:

91. 92
•The nonbonding electrons in an atom are called
lone pairs.
•There can be an expanded octet if it is necessary to
increase the number of electrons around a central
atom. In this case, the d orbital electrons are used in
bonding. Two or four electrons may be added. The
formal charges are used to find out if the expanded
octet is more stable than the Lewis electron dot
structure.
•There are some exceptions of the octet rule when
making the Lewis electron dot structure of
compounds. There are some elements that can be
stable by having less than 8 valence electrons.

92. 93
Complex Lewis StructuresComplex Lewis Structures

93. 94
There are some molecules and
polyatomic ions for which no single
Lewis structure consistent with all
characteristics and bonding information
can be written.

94. 95
Step 1. The total number of valence electrons
is 24, 5 from the nitrogen atom and 6 from
each O atom, and 1 from the –1 charge.
Write a Lewis structure for NO
2.
-
3
NO.

95. 96
Since the extra electron present results in
nitrate having a –1 charge, the ion is enclosed
in brackets with a – charge.
Step 2. The three O atoms are bonded to a
central N atom. Write the skeletal structure
and place two electrons between each pair of
atoms.
Write a Lewis structure for NO
2.
-
3
NO.
O N O::
O:
-

96. 97
Step 3. Subtract the 6 electrons used in Step 2
from 24, the total number of valence electrons,
to obtain 18 electrons yet to be placed.
Since the extra electron present results in
nitrate having a –1 charge, the ion is enclosed
in brackets with a – charge.
O N O::
O:
-
Write a Lewis structure for NO
2.
-
3
NO.

97. 98
O N O
O
Step 4. Distribute the 18 electrons around the
N and O atoms.
Write a Lewis structure for NO
2.
-
3
NO.
:
:::
::
:: :
::
:
electron deficient

98. 99
Write a Lewis structure for NO
2.
-
3
NO.
Step 5. One pair of electrons is still needed to
give all the N and O atoms a noble gas
structure. Move the unbonded pair of electrons
from the N atom and place it between the N
and the electron-deficient O atom, making a
double bond.
:
:::
::
:: :
O N O
O
::
:
-

99. 100
Write a Lewis structure for NO
2.
-
3
NO.
Step 5. One pair of electrons is still needed to
give all the N and O atoms a noble gas
structure. Move the unbonded pair of electrons
from the N atom and place it between the N
and the electron-deficient O atom, making a
double bond.
N
O
:
:
O:
:
:
O:
:
:
:
-

100. 101
A molecule or ion that shows multiple
correct Lewis structures exhibits resonance.
Write a Lewis structure for NO
2.
-
3
NO.
Step 5. There are three possible Lewis
structures.
N
O
:
:
O:
:
:
O:
:
:
:
-
:
N
O
:
:
O:
:
O:
:
:
:
-
Each Lewis structure is called a resonance
structure.
N
O
:
:
O:
:
:
O:
:
:
:
-

101. 102
A molecule or ion that shows multiple
correct Lewis structures exhibits resonance.
Write a Lewis structure for NO
2.
-
3
NO.
Step 5. There are three possible Lewis
structures.
N
O
:
:
O:
:
:
O:
:
:
:
-
:
N
O
:
:
O:
:
O:
:
:
:
-
Each Lewis structure is called a resonance
structure.
N
O
:
:
O:
:
:
O:
:
:
:
-

102. 103
Compounds ContainingCompounds Containing
Polyatomic IonsPolyatomic Ions

103. 104
A polyatomic ion is a stable group of
atoms that has either a positive or
negative charge and behaves as a single
unit in many chemical reactions.

104. 105
Sodium nitrate, NaNO
3, contains one
sodium ion and one nitrate ion.
sodium ion Na
+
-
3
NOnitrate ion
N
O
:
:
O:
:
:
O:
:
:
:
-
Na
+

105. 106
•The nitrate ion is a polyatomic ion
composed of one nitrogen atom and
three oxygen atoms.
N
O
:
:
O:
:
:
O:
:
:
:
-
Na
+
•It has a charge of –1
•One nitrogen and three oxygen atoms
have a total of 23 valence electrons.

106. 107
•The –1 charge on nitrate adds an
additional valence electron for a total of
24.
N
O
:
:
O:
:
:
O:
:
:
:
-
Na
+
•The additional valence electron comes
from a sodium atom which becomes a
sodium ion.

107. 108
•Sodium nitrate has both ionic and
covalent bonds.
N
O
:
:
O:
:
:
O:
:
:
:
-
Na
+
•Ionic bonds exist between the sodium
ions and the carbonate ions.
covalent
bond
covalent
bond
covalent
bond
ionic
bond
•Covalent bonds are present between the
nitrogen and oxygen atoms within the
nitrate ion.

108. 109
•When sodium nitrate is dissolved in
water, the ionic bond breaks.
N
O
:
:
O:
:
:
O:
:
:
:
-
Na
+
•The sodium ions and nitrate ions separate
from each other forming separate sodium
and nitrate ions.
N
O
:
:
O:
:
:
O:
:
:
:
-
Na
+
•The nitrate ion, which is held together by
covalent bonds, remains as a unit.

109. 110
FORMAL CHARGE

110. 111
Formal charge is the charge of the
atom based on the Lewis electron dot
structure for a particular compound.
The formal charge can be calculated using the
following:
Formal Charge (FC) = group number – (number
of unshared electrons + number of bonds)

111. 112
•Example:
A compound made up of H, O, and Br. There are
two possible structures with either Br or O as the
central atom.
Formal Charges for H – Br – O
For H: FC = 1 – (0 + 1) = 0
For Br: FC = 7 – (4 + 2) = 1
For O: FC = 6 – (6 + 1) = -1
Formal Charge for H – O – Br
For H: FC = 1 – (0 + 1) = 0
For Br: FC = 7 – (6 + 1) = 0
For O: FC = 6 – (4 + 2) = 0
 
Those with 0 formal charges are the most possible structure.

112. 113
Molecular ShapeMolecular Shape

113. 114
The 3-dimensional arrangement of the
atoms within a molecule is a significant
determinant of molecular interactions.

114. 115
11.12

115. 116
The Valence ShellThe Valence Shell
Electron Pair (VSEPR) Electron Pair (VSEPR)
ModelModel

116. 117
The VSEPR model is based on the idea
that electron pairs will repel each other
electrically and will seek to minimize
this repulsion.
To accomplish this minimization, the
electron pairs will be arranged as far
apart as possible around a central atom.

117. 118
BeCl
2 is a molecule with only two pairs of
electrons around beryllium, its central
atom.
Its electrons are arranged 180
o
apart for
maximum separation.

118. 119
•BF
3 is a molecule with three pairs of electrons
around boron, its central atom.
•Its electrons are arranged 120
o
apart for
maximum separation.
•This arrangement of atoms is called trigonal
planar.

119. 120
•CH
4 is a molecule with four pairs of electrons
around carbon, its central atom.
•An obvious choice for its atomic
arrangement is a 90
o
angle between its atoms
with all of its atoms in a single plane.
•However, since the molecule is 3-dimensional
the molecular structure is tetrahedral with a bond
angle of 109.5
o
.

120. 121
Ball and stick models of methane, CH
4, and carbon
tetrachloride, CCl
4.
11.13

121. 122
•Ammonia, NH
3, has four electron pairs
around nitrogen.
The arrangement
of electron pairs
around nitrogen is
tetrahedral.

122. 123
The NH
3
molecule
is pyramidal.
NH
3
has one
unbonded pair
of electrons.

123. 124
•Water has four electron pairs around
oxygen.
The arrangement
of electron pairs
around oxygen is
tetrahedral.

124. 125
The H
2O
molecule
is bent.
H
2
O

has two
unbonded pairs
of electrons.

125. 126