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1
Chapter 7
Chemical Bonding
2
Nature of Chemical Bond• Atoms are held together by electrostatic attraction between
positively charged nuclei and negatively charged electron clouds.
• Chemical Bond: a link between atoms that result from mutual attraction of their nuclei for electrons.
• Bond energy: the energy required to break a bond.
• Forces in substances:– Attractive: between electron clouds and respective bonding
nuclei (of the two atoms that bond)– Repulsive: between all the electron clouds in the bonding
atoms
3
Formation of Bonds• Bonding involves only the valence electrons (those in
the highest energy level).
• Use the periodic chart to guide determination of valence electrons
• WHEN BONDING OCCURS:
– Atoms attain an OCTET: a stable Noble Gas configuration.
– the resulting system is at the lowest possible potential energy level.
– The process of bonding is, therefore, exothermic: energy is being released. If the energy released is Large we get a strong bond; small ΔE bond is weak
4
Types of Bonds
• Ionic bond: formed by transfer of electrons from the valence energy level of one atom to another’s
• Covalent bond: formed when atoms share electrons.
• Metallic bond: ions of metals are surrounded by sea of electrons that bind all ions together.
5
Types of Bonds: Ionic
• Ionic Bond: results from electrostatic attraction between positive and negative ions.
• Produced by TRANSFER of electrons from valence energy level of one atom to another
• Occurs between metal and nonmetals , or polyatomic ions.
6
Types of Covalent Bonding
Types
• Nonpolar
• Polar
• Coordinate Covalent Bond
• Network covalent
7
Electronegativity
• A measure of how strongly the atoms attract electrons in a bond.
• The bigger the electronegativity difference the more polar the bond.
• 0.0 - 0.3 Covalent nonpolar• 0.3 - 1.0 Covalent moderately polar• 1.0 -1.7 Covalent polar• >1.7 Ionic• Use table 6.6, page 171 in your textbook
8
Pg 335
Table 8-1Representative Electronegativity
Differences
Covalent: = 0
Polar:0.3 < < 1.7
Ionic > 1.7
9
Nonpolar Covalent Bonding
• Nonpolar covalent bond: electrons are shared equally by atoms.– Electronegativity difference <0.3
• Examples:H-H O2
Cl-Cl Si-HN2 Ge-H
10
Polar Covalent Bonding
• Polar covalent bonds: the electrons are not shared equally between the bonding atoms. The more electronegative atom attracts the electrons.– Electronegativity difference between >0.3
and <1.7• One end is slightly positive, the other negative. • The charge distribution is indicated using small
delta and
11
A + indicates that the atom that doesn’t hog the electron density (less electronegative) has a partial positive charge. A - indicates the atom that hogs the electrons in the molecule (is more electronegative)
H F FH
Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms
electron richregion
electron poorregion e- riche- poor
+ -
9.5
12
H - F+ -
H - F
+-H - F+
-
H - F
+-
H - F +-
H - F+-
H - F
+-
H - F
+-
13
H - F+ -
H - F
+-H - F+
-
H - F
+-
H - F +-
H - F+-
H - F
+-
H - F
+-
+-
14
H - F+ -
H - F+ -
H - F+ - H - F
+ -
H - F+ -
H - F+ -
H - F+ -
H - F+ -
- +
15
Examples
• Determine the nature of the bond (nonpolar covalent, polar covalent, or ionic) for the following two elements:
1. Na and Cl 6. N and O
2. Si and F 7. S and F
3. S and Br 8. S and H
4. C and H 9. Te and I
5. Al and F 10. K and I
16
Network Covalent Bond
• Occurs in compounds (elements) where all the atoms are bonded with covalent bonds .
• Examples: diamond, graphite
diamond
17
AllotropesDifferent Structural forms of the same
element
Carbon BuckyballsAmorphous
Graphite
18
Coordinate Covalent Bond
• Additional Examples:– NH3 + H+ → NH4
+
– AlCl3 + Cl-1 → AlCl4-1
– Al(OH)3 + OH- → Al(OH)4-1
19
Lewis Theory: An Overview
• Valence e- play a fundamental role in chemical bonding.
• e- transfer leads to ionic bonds.
• Sharing of e- leads to covalent bonds.
• e- are transferred or shared to give each atom a noble gas configuration – the octet.
20
Lewis Symbols (Structures)
• A chemical symbol represents the nucleus and the core e-.
• Dots around the symbol represent valence e-.
Si•
••
•
N••
••
• P••
••
• As••
••
• Sb••
••
• Bi••
••
•
••Al••
• Se••
•••
Ar••
••
••I •••
••
••
21
Ionic and Molecular Compounds
• Formation of sodium chloride (ionic):
• Formation of hydrogen chloride (covalent):
A metal and a nonmetal transfer electrons to form an ionic compound. Two nonmetals share electrons to form a molecular compound.
Na + Na+ [ ]Cl
Cl
H + Cl
Cl
H
22
Ionic Compounds
Ionic compounds consist of
a lattice
of positive
and negative ions.
Lattice: three dimensional array of ions
NaCl:
23
Properties of Ions
• When will a stable bond be formed?• Octet Rule: both ions attain noble gas
configuration
• Example: NaCl versus Na+Cl-
Na: [Ne]3s1 Cl: [Ne]3s23p5
Na+: [Ne] Cl-: [Ne]3s23p6 = [Ar]
STRONG ELECTROSTATIC ATTRACTION
24
Ionic Bonding Formation(1)• An atom with a low ionization energy reacts with
an atom with high electron affinity. • Between active metal (positive ion) and active
nonmetal (negative ion), or between polyatomic ions.
• Opposite charges hold the atoms together.
• Electronegativity difference – Electronegativity difference between atoms
>1.7 – 50% or more ionic character (check PT)
25
Ionic Bond Formation(2)
• Both the positive and negative ions acquire noble gas configuration.
• The charge of the ion (oxidation state) is determined by the number of electrons lost or gained.
• The grater the electronegativity difference between the elements, the greater the ionic character of the bond. (>1.7 , 50% ionic).
• Only valence electrons participate in bond formation ( exception: transition elements).
26
Lewis Structures for Ionic Compounds
• CHECK THE BLACKBOARD (Lewis structures and Orbital Notations)
• NaCl• MgCl2• Al2O3
• MgO• AlF3
27
Properties of Ionic Compounds1. Crystalline structure.
A regular repeating arrangement of ions in the solid (lattice).
2. Ions are strongly bonded.3. Structure is rigid.4. High melting points
due to strong forces between ions.5. Oxidation states (charges) are determined by the #
of valence electrons (Group #)6. Conduct electricity when molten and in aqueous
solution7. Do not reflect light – therefore are white
(exceptions: transition elements ions)8. Solubility in water depends on lattice energy and
nature of solvent.
28
Crystalline structure
29
Ionic solids are brittle
+ - + -+- +-
+ - + -+- +-
30
Ionic solids are brittle
+ - + -
+- +-+ - + -
+- +-
• Strong Repulsion breaks crystal apart.
31
Metallic Bond
32
Metallic Bond
• Metals consist of crystalline lattice in which positive ions (kernels) are arranged in fixed patterns.
• The valence electrons are free to move and they belong to the entire crystal.
• “Electron Sea” model
33
Metallic Bond
34
COVALENT BONDS
35
What about covalent compounds?
The electrons in each atom are attracted to the nucleus of the other.
The electrons repel each other,The nuclei repel each other.
The atoms reach a distance with the lowest possible energy.
The distance between the atoms is the bond length.
36
When Atoms Combine to make Molecules
Fig 8-1
Atoms contain both positive and negative charges. When they come Together they arrange themselves so that the attractive forges of oppositeCharges is greater than the repulsive forces of like charges
37
Fig 8-3 Pg 330
The interaction energy of a pair of hydrogen atoms varies with internuclear separation.
38
How does H2 form?
• The nuclei repel
++
39
How does H2 form?
• The nuclei repel
• But they are attracted to electrons
• They share the electrons
++
40
Covalent Bond Formation
Covalent bond forms by overlap of orbitals.• Two types of bonds
Sigma bond: all single bonds are sigma bonds ( along the internuclear
plane)
Pi bond: in multiple bonds: the first one is sigma, all other bonds are pi.( above/below and front/back)
There areSingle bondsMultiple bonds (double and triple only)
41
42
Two types of Bonds
• Sigma bonds(σ) from overlap of orbitals along the axis connecting the nuclei between the atoms
• Pi bond (): perpendicular overlap of p-orbitals above and below the axis connecting the atoms
43
Sigma Bond Formation
• s-s overlap: Overlap of two s orbitals (s-s) : H-H bond
• p-p overlap: Overlap of two p orbitals (p-p) facing each other along the same axis (x-axis): Cl-Cl bond
• s-p overlap: overlap of s-orbital and p-orbital along the same axis. H-Cl bond
44
Sigma Bond Formation: s-s Orbital Overlap
2 single atoms start overlap
Overlap complete
45
Sigma Bond Formation p-p overlap
• P-orbital overlap
46
Sigma bond: p – p overlap
47
Pi Bond
• Forms by a vertical overlap of two p orbitals (p-p vertical overlap)
• Exists only when there are multiple covalent bonds
• Example H2 -C=C-H2: the double
bonds contains one sigma and one pi bond
48
Sideways overlap of p-orbitals to form a pi - bond
49
Pi Bond
50
51
Pi-bond formation in ethene.
52
Movies of Bond Formation
• http://www.google.com/imgres?imgurl=http://www.cem.msu.edu/~harrison/johnston/movies/nitrogen/sigma7icon.jpg&imgrefurl=http://www.cem.msu.edu/~harrison/johnston/nitrogen.html&h=92&w=99&sz=5&tbnid=aJAk-oo0W7Oe7M:&tbnh=71&tbnw=77&hl=en&start=11&prev=/images%
• http://www.google.com/imgres?imgurl=http://www.cem.msu.edu/~harrison/johnston/movies/nitrogen/sigma7icon.jpg&imgrefurl=http://www.cem.msu.edu/~harrison/johnston/nitrogen.html&h=92&w=99&sz=5&tbnid=aJAk-oo0W7Oe7M:&tbnh=71&tbnw=77&hl=en&start=11&prev=/images%
53
Depiction of Covalent Compounds
• Covalent compounds can be described by
– Molecular formula (molecular compounds)
– Structural formula (depicts the arrangement of the atoms in space)
– Lewis structure: depicts the arrangement of the electrons around the atoms in a molecule
54
Writing Lewis Structures
• The Lewis Structures will be written for molecules that obey the Octet Rule
• AND
• Exceptions to the Octet Rule– Electron deficiency– Expanded Octet Rule– Odd electron molecules
55
Writing Lewis Structures
Blackboard
56
Resonance
• When more than one dot diagram with the same connections are possible.
• Use double arrows to indicate it is the “average” of the structures.
• NO2-
• Which one is it?• Does it go back and forth.• It is a mixture of both, like a mule.• NO3
- CO3-2 SO3
57
Molecular Geometry
58
Molecular Geometry
• Lewis structures tell us how the atoms are connected to each other.
• They don’t tell us anything about shape.
• The shape of a molecule can greatly affect its properties.
59
Molecular Geometry
Three theories to explain:
• VSEPR: Valence shell electron pair repulsion theory.
• Valence Bond Theory: Hybridization theory
• Molecular Orbital Theory: advanced course
60
Importance of Molecular Shape
Three dimensional structure of a molecule can have a profound effect on its reactivity and biological activity.
These two molecules have identical formulas and shape, but they are mirror images of each other and they have different pharmacological activity (PA) .
Enantiomers and optical isomers
61
Hemoglobin
62
caffeine
penicillin
AmmoniaWater
63
VSEPR• Lewis structures tell us how the atoms are
connected to each other.
• They don’t tell us anything about shape.
• The shape of a molecule can greatly affect its properties.
• Valence Shell Electron Pair Repulsion Theory allows us to predict geometry
64
VSEPR Theory
• Electron Pair Repulsion Theory: electron pairs (both shared and unshared) try to orient themselves as far away as possible in the space around the central atom.
65
VSEPR Theory• Uses Lewis structures and shared and
unshared pair of electrons to predict geometry.
• All electrons are in their original atomic orbitals.
• Predicts three dimensional geometry of molecules. Can predict the angles of bonds.
66
Properties of Shared and Unshared Electron Pairs(1)
• Shared pair of electron: the electron pair is attracted by both nuclei – shape slender, like a cigar.
• Unshared pairs: take a lot of space as electrons repulse each other. Shaped like a pear, or Mickey Mouse ears.
67
Properties of Electron Pairs (2)
• The strength of repulsions between pair of electrons:
unshared-unshared > shared - unshared>
shared - shared • Molecular shape: repulsions between
charge cloud determines the arrangement
68
VSEPR• Molecules take a shape that puts electron
pairs as far away from each other as possible.
• Draw the Lewis structure to determine electron pairs.
• Determine:–bonding–nonbonding lone pair
• Lone pair take more space.• Multiple bonds count as one pair.
69
VSEPR
• The number of pairs determines
–bond angles
–underlying structure
• The number of atoms determines
–actual shape
70
Electron-Group Geometries
• Differentiate between electron-group geometries and molecular geometry
• 2 electron groups: linear
• 3 electron groups: trigonal planar
• 4 electron groups: tetrahedral
• 5 electron groups: trigonal bipyramidal
• 6 electron groups: octahedral
71http://www.mpcfaculty.net/mark_bishop/tri-plan.htm
Electron Group Geometries
72
Electron Group Geometries
Electronpairs
BondAngles
UnderlyingShape
2 180° Linear
3 120° Trigonal Planar
4 109.5° Tetrahedral
590° &120°
Trigonal Bipyramidal
6 90° Octagonal
73
Valence shell electron pair repulsion (VSEPR) model:N0 LONE Electron Pairs around the Central Atom
Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.
AB22 0
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
10.1
linear linear
B B
74
Two electron-group geometries
75
Molecular Geometry
No Lone Pairs of Electrons on Central Atom
76
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: all pairs are shared (0 lonely pairs)
AB3 3 0trigonal planar
trigonal planar
10.1
7710.1
3 electron-group geometry
Other examples: SO3
78
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: all pairs are shared
AB3 3 0trigonal planar
trigonal planar
10.1
AB4 4 0 tetrahedral tetrahedral
79
Tetrahedral Structure
8010.1
4 electron-group geometry
Other examples: NH4+1, SO4
-2
81
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: All electron pairs are shared
AB3 3 0trigonal planar
trigonal planar
10.1
AB4 4 0 tetrahedral tetrahedral
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
8210.1
5 electron-group geometry
83
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: all electron pairs are shared
AB3 3 0trigonal planar
trigonal planar
10.1
AB4 4 0 tetrahedral tetrahedral
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB6 6 0 octahedraloctahedral
8410.1
6 electron-pair geometry
85
Tetrahedral Structure
8610.1
5 electron-group geometry
8710.1
6 electron-pair geometry
88
Molecular Geometry
Lone Electron Pairs on Central Atom
89
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: THREE ELECTRON-GROUPS
AB3 3 0trigonal planar
trigonal planar
AB2E 2 1 same bent
10.1
90
91
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: FOUR ELECTRON-GROUPS
AB3E 3 1
AB4 4 0 tetrahedral tetrahedral
tetrahedraltrigonal
pyramidal
10.1
92
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: FOUR ELECTRON-GROUPS
AB4 4 0 tetrahedral tetrahedral
10.1
AB3E 3 1 tetrahedraltrigonal
pyramidal
AB2E2 2 2 tetrahedral bent
H
O
H
93
94bonding-pair vs. bonding
pair repulsionlone-pair vs. lone pair
repulsionlone-pair vs. bonding
pair repulsion> >
Effect of Unshared Electron Pair on Bond Angles
95
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: FIVE ELECTRON-GROUPS
10.1
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
See-saw
96
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: FIVE ELECTRON-GROUPS
10.1
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
AB3E2 3 2trigonal
bipyramidalT-shaped
ClF
F
F
97
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: FIVE ELECTRON-GROUPS ( LONE PAIRS)
10.1
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
AB3E2 3 2trigonal
bipyramidalT-shaped
AB2E3 2 3trigonal
bipyramidallinear
I
I
I
98
Five Electron Pairs on Central Atom
99
100
101
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: SIX ELECTRON-GROUPS ( LONE PAIRS)
10.1
AB6 6 0 octahedraloctahedral
AB5E 5 1 octahedral square pyramidal
Br
F F
FF
F
102
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR: SIX ELECTRON-GROUPS ( LONE PAIRS)
10.1
AB6 6 0 octahedraloctahedral
AB5E 5 1 octahedral square pyramidal
AB4E2 4 2 octahedral square planar
Xe
F F
FF
103
104
105
106
107
Angles in the Different Geometries
108
Angles in the Different Geometries
109
VSEPRElectron
pairsBond
AnglesUnderlyingShape
2 180° Linear
3 120° Trigonal Planar
4 109.5° Tetrahedral
590° &120°
Trigonal Bipyramidal
6 90° Octagonal
110
Actual shape
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
2 2 0 linear
3 3 0 trigonal planar
3 2 1 bent4 4 0 tetrahedral4 3 1 trigonal pyramidal4 2 2 bent
111
Actual Shape
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
5 5 0 trigonal bipyrimidal
5 4 1 See-saw
5 3 2 T-shaped5 2 3 linear
112
Actual Shape
ElectronPairs
BondingPairs
Non-Bonding
Pairs Shape
6 6 0 Octahedral
6 5 1 Square Pyramidal
6 4 2 Square Planar6 3 3 T-shaped6 2 1 linear
113
VSEPR: Examples
• CH4 - draw the structural formula
• Determine the number of shared and unshared pairs of electrons.
• Write the Lewis structure
114
VSEPR
• Single bonds fill all atoms.
• There are 4 pairs of electrons pushing away.
• The furthest they can get away is 109.5º (REMEMBER: three dimensional space.
C HH
H
H
115
4 atoms bonded
• Basic shape is tetrahedral.
• A pyramid with a triangular base.
• Same shape for everything with 4 pairs.
CH HH
H109.5º
116
3 bonded - 1 lone pair
N HH
H
NH HH
<109.5º
• Still basic tetrahedral but you can’t see the electron pair.
• Shape is calledtrigonal pyramidal.
117
2 bonded - 2 lone pair
OH
H
O HH
<109.5º
• Still basic tetrahedral but you can’t see the 2 lone pair.
• Shape is calledbent.
118
Geometry Using VSEPR Theory
• Vocabulary: – A: central atom– B: bonded atom– E: unshared pair of electrons
• CHECK BLACKBOARD FOR TYPE OF GEOMETRY AND EXAMPLES.
119
Linear Geometry ( 3 atoms)
AB B
Designation: AB2
120
Three Electron Densities
A
B
B
B
B
B
A
E
AB3, trigonal planar AB3E, Bent
121
Four Electron Densities
AB4 AB3E AB2E2
122
Five Electron Densities
AB5
AB3E2
AB4E
AB2E3
123
Six Electron Densities
AB6 AB5E
AB4 E2
124
3 Atoms No Lone Pair (double Bond)
CH
HO
• The farthest you can the electron pair apart is 120º
125
3 atoms no lone pair
CH
HO
• The farthest you can place the electron pairs apart is 120º.
• Shape is flat and called trigonal planar.
C
H
H O
120º
126
2 atoms no lone pair
• With three atoms the farthest they can get apart is 180º.
• Shape called linear.
C OO180º
127
Valence Bond Theory: Hybrid Orbitals
Combines bonding with geometry
128
Hybridization
• The mixing of several atomic orbitals to form the same number of hybrid orbitals.
• All the hybrid orbitals that form are the same.
129
Types of Hybrid Orbitals
• sp3 : 1 s and 3 p orbitals mix to form 4 sp3 orbitals.
• sp2 :1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital intact.
• sp : s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals intact.
• sp3d: five orbitals• sp3d2 : six orbitals
130
Hybridization
• We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry.
• We combine one s orbital and 3 p orbitals.
• sp3 hybridization has tetrahedral geometry.
131
How we get to hybridization
• We know the geometry from experiment.
• We know the orbitals of the atom
• hybridizing atomic orbitals can explain the geometry.
• So if the geometry requires a tetrahedral shape, it is sp3 hybridized.
• This includes bent and trigonal pyramidal molecules because one of the sp3 lobes holds the lone pair.
132
Hybridization
• We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry.
• We combine one s orbital and 3 p orbitals.
• sp3 hybridization has tetrahedral geometry.
133
sp3 geometry
109.5º
• This leads to tetrahedral shape.
• Every molecule with a total of 4 atoms and lone pair is sp3 hybridized.
• Gives us trigonal pyramidal and bent shapes also.
134
135
136
sp2 hybridization
• C2H4, BF3
• double bond ( in C2H4 ) acts as one pair
• trigonal planar
• Have to end up with three blended orbitals
• use one s and two p orbitals to make sp2 orbitals.
• leaves one p orbital perpendicular
137
138
139
140
sp3 geometry
109.5º
• This leads to tetrahedral shape.
• Every molecule with a total of 4 atoms and lone pair is sp3 hybridized.
• Gives us trigonal pyramidal and bent shapes also.
141
sp2 hybridization in C2H4
• trigonal planar
• 120º angle
• one bond
• One sigma and one pi bond between the C-C atoms
142
Where is the P orbital?
• Perpendicular
• The overlap of orbitals makes a sigma bond ( bond)
143
CCH
H
H
H
144
What about sp
• one s and one p hybridize
• Linear
• C2H2
145
sp hybridization
• end up with two lobes 180º apart.
• p orbitals are at right angles
• makes room for two bonds and two sigma bonds.
• a triple bond or two double bonds
146
Hybridization and Geometry of Electrons
• sp3: tetrahedral• sp2: trigonal planar• sp: linear• sp3d: trigonal bipyramid• sp3d2: octahedral• Geometry of molecule determined by
number of shared and unshared electrons.
147
Types of Hybrid Orbitals
• sp3 : 1 s and 3 p orbitals mix to form 4 sp3 orbitals.
• sp2 :1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital intact.
• sp : s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals intact.
• sp3d: five orbitals• sp3d2 : six orbitals
148
Polar Bonds
• When the atoms in a bond are the same, the electrons are shared equally.
• This is a nonpolar covalent bond.
• When two different atoms are connected, the electrons may not be shared equally.
• This is a polar covalent bond.
• How do we measure how strong the atoms pull on electrons?
149
How to show a bond is polar• Isn’t a whole charge just a partial charge means a partially positive means a partially negative
• The Cl pulls harder on the electrons
• The electrons spend more time near the Cl
H Cl
150
Hybridization and Geometry of Electrons
• sp3: tetrahedral• sp2: trigonal planar• sp: linear• sp3d: trigonal bipyramid• sp3d2: octahedral• Geometry of molecule determined by
number of shared and unshared electrons.
151
Types of Hybrid Orbitals
• sp3 : 1 s and 3 p orbitals mix to form 4 sp3 orbitals.
• sp2 :1 s and 2 p orbitals mix to form 3 sp2 orbitals leaving 1 p orbital intact.
• sp : s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals intact.
• sp3d: five orbitals• sp3d2 : six orbitals
152
Polar Bonds
• When the atoms in a bond are the same, the electrons are shared equally.
• This is a nonpolar covalent bond.
• When two different atoms are connected, the atoms may not be shared equally.
• This is a polar covalent bond.
• How do we measure how strong the atoms pull on electrons?
153
How to show a bond is polar• Isn’t a whole charge just a partial charge means a partially positive means a partially negative
• The Cl pulls harder on the electrons
• The electrons spend more time near the Cl
H Cl
154
Polar Molecules
Molecules with endsDetermined by polarity of bonds
AndSymmetry of Molecules
155
Partial Ionic Compounds (cont.)
Covalent
Polar Covalent
Ionic
Increased Ionic Character
156
Criteria for Polarity of molecules
• Requires two things to be true The molecule must contain polar bonds. This can be determined from differences in
electronegativity.Symmetry can not cancel out the effects
of the polar bonds. Must determine geometry first.
157
Criteria for Polarity of molecules
• Requires two things to be true The molecule must contain polar bonds. This can be determined from differences in
electronegativity.Symmetry can not cancel out the effects
of the polar bonds. Must determine geometry first.
158
Geometry and polarity• Three shapes will cancel them out.
• Linear
159
Geometry and polarity• Three shapes will cancel them out.
• Planar triangles
120º
160
Geometry and polarity• Three shapes will cancel them out.
• Tetrahedral
161
Geometry and polarity• Others don’t cancel
• Bent
162
Geometry and polarity• Others don’t cancel
• Trigonal Pyramidal
163
Is it polar?
• HF
• H2O
• NH3
• CCl4
• CO2
• PCl5• CO
164
Bond Dissociation Energy
• The energy required to break a bond
• C - H + 393 kJ C + H
• We get the Bond dissociation energy back when the atoms are put back together
• If we add up the BDE of the reactants and subtract the BDE of the products we can determine the energy of the reaction (H)
165
Find the energy change for the reaction
• CH4 + 2O2 CO2 + 2H2O
• For the reactants we need to break 4 C-H bonds at 393 kJ/mol and 2 O=O bonds at 495 kJ/mol= 2562 kJ/mol
• For the products we form 2 C=O at 736 kJ/mol and 4 O-H bonds at 464 kJ/mol
• = 3328 kJ/mol
• reactants - products = 2562-3328 = -766kJ