Title: Unit 11: Chemical Bonding
1Unit 11 Chemical Bonding
2Chemical Bonding Overview
- Chemical Bond
-
- attractive force is between the nucleus of one
atom and the valence electrons of another atom.
3Review
- How do you find valence electrons?
- Hint there are two ways!
- Examples
- Mg ___
- O ___
- Ar ___
- Si ___
- Examples
- Mg
- O
2
6
8
4
1s2 2s2 2p6 3s2
1s2 2s2 2p4
4Electron Dot Structures
- Depicts element symbol w/ valence e- shown as
dots.
Na
Mg
Al
Si
Cl
Ar
O
N
5Purpose of Bonding
- Chemical bonding is driven by the octet rule
atoms will lose, gain, or share electrons in
order to achieve the electron configuration of
the closest noble gas. - Unfilled or partially filled valence orbitals are
inherently unstable. (Unstable means possessing
high potential energy.) Nature always strives for
the lowest energy conditions
6Types of Chemical Bonds
- There are 3 types of chemical bonds
- Ionic Covalent Metallic
- Depends on the number of valence electrons of
bonding atoms. - Metallic bonds form between metal atoms only.
- Ionic bonds involve the transfer of electrons
(lose/gain). - Covalent bonds involve the sharing of electrons.
- Metallic bonds involve a mobile electron sea.
Flower Time
7Movie
8Ionic Bonds
- Occurs when ions of opposite charge (,-) attract
each other. - Metal ion Nonmetal ion
- Simplest attraction
- NaCl MgF2
- Polyatomic ions
- AlPO4 (NH4)2SO4
9Formation of Ionic Bond
- Cation- positive ion ()
- Forms when a metal atom loses e- to become
stable. - Anion- negative ion (-)
- Forms when a nonmetal atom gains e- to become
stable - An ionic bond is formed when e- are transferred
between atoms and the resulting ions stick
together.
10Examples
- Formation of NaCl
- Na Cl ? Na Cl - NaCl
- Formation of MgF2
- Mg F F ? Mg2 F - F - MgF2
- How would aluminum oxide form?
11Electron Configuration of Ions
- Cation example (metal)
- Ca atom 1s2 2s2 2p6 3s2 3p6 4s2
- Ca2 ion 1s2 2s2 2p6 3s2 3p6 Lost 2 electrons
to obtain noble gas configuration (octet)
12Electron Configuration of Ions
- Anion example (nonmetal)
- N atom 1s2 2s2 2p3
- N3- ion 1s2 2s2 2p6
- Gained 3 electrons to obtain noble gas
configuration (octet)
13Properties of Ionic Compounds
IONIC
e- transferred from metal to nonmetal
Bond Formation
Type of Structure
Crystal lattice
Physical State
Solid (hard and rigid)
Melting Point
high
Boiling Point
high
yes (solution or liquid)
Electrical Conductivity
Other Properties
brittle
14Electrolyte
- A substance that conducts electricity
- Because of ionic bonds ionic (charged) nature,
ionic compounds conduct electricity in the molten
or aqueous forms.
15Bonding Types
16Bond Polarity
17NON-POLAR COVALENT BONDS
18Polar Covalent Bonds Unevenly matched, but
willing to share.
19Ionic Bonds One Big Greedy Thief Dog!
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21Bond Polarity
- Nonpolar Covalent Bond
- e- are shared equally
- symmetrical e- density
- usually identical atoms
22Bond Polarity
- Polar Covalent Bond
- e- are shared unequally
- asymmetrical e- density
- results in partial charges (dipole)
23Bond Polarity
- Most bonds are a blend of ionic and covalent
characteristics. - Difference in electronegativity determines bond
type.
If DEN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt DEN lt 1.7 Polar covalent
gt 1.7 Ionic
24- A large electronegativity difference leads to an
ionic bond. - A small electronegativity difference leads to a
polar covalent bond. - No electronegativity difference between two atoms
leads to a nonpolar covalent bond.
25- Cesium and sulfur
- Carbon and oxygen
- Fluorine and fluorine
If ?EN is Bond type is
lt 0.4 Nonpolar covalent
0.4 lt ? EN lt 1.7 Polar covalent
gt 1.7 Ionic
Ionic Polar
Nonpolar
26Covalent Bonds
- Occurs when 2 nonmetals share pairs of electrons
to become stable. Molecular compounds are
formed. - Examples
- H2O CO2 C6H12O6 PCl5
27Covalent Bonds
- Covalent bonds can be
- single (1 shared pair)
- double (2 shared pairs)
- or triple (3 shared pairs)
- Bond strength triple gt double gt single
- Bond length single gt double gt triple
28Lewis Structures
- Creating Lewis Structures
- Lewis structures are depictions of molecules that
show valence electrons as __dots________. - Shared pairs of electrons (i.e. _Valence
electrons_) are drawn between the atoms sharing
them. - Unshared or ___lone_______ pairs of electrons are
represented by dots located on one atom only
298
2
NEED
4
8
8
8
8
6
8
308
1
Available
2
4
5
6
7
3
Figure out through e config
31Creating Lewis Structures
- Follow this system
- Example H2O
- 1) Draw a skeleton of the molecule. It
generally works to place the different atom in
the center. - H O H
32Creating Lewis Structures
- Find the needed electrons (N) for each atom and
add them up. N will be 8 for most elements, with
these exceptions
- H gets 2 valence e-
- Be gets 4 valence e-
- B gets 6 valence e-
33- 3) Find the available (valence) electrons (A)
for each atom and then add them up. - A
- special note when completing a Lewis structure
for a polyatomic ion, you will need to correct A
by adding the absolute value of the charge if
negative, and subtracting the charge if positive.
For example, for the ion PO43-, you would add 3
to A. For the ion NH4, you would subtract 1
from A. (You do the opposite of the charge.)
H 1 O 6 H 1 Total A 8
N 12 A 8
34- 4) Find the shared (S) electrons for the entire
molecule by this formula S N A - S
S 12 8 4
N 12 A 8 S 4
35- 5) The shared electrons are the bonding
electrons. Place all of the shared electrons
between the atoms. - H O H
- 6) You must place all of the available (A)
electrons in the picture. The shared electrons
are part of the available. See how many of the
available electrons still need to be placed, and
put them in the picture as lone pairs (unshared
pairs) so that every atom gets an octet (remember
H only needs 2). - H O H
N 12 A 8 S 4
N 12 A 8 S 4
4
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39H H C H H
N A S
8(2x4) 16
4(1x4) 8
0
8
40F F C F F
N A S
8(4x8) 40
4(4x7) 32
24
8
41Cl Cl P Cl
N A S
8(3x8) 32
5(3x7) 26
20
6
42 O C O
N A S
8(2x8) 24
4(2x6) 16
8
8
43Polyatomic Ions
- To find total of valence e- (A)
- Add 1e- for each negative charge.
- Subtract 1e- for each positive charge.
- Place brackets around the ion and label the
charge.
44Polyatomic Ions
N A S
8(4x8) 40
O O Cl O O
7(4x6) 31
1 32
24
8
45Resonance
- Resonance structures are structures that occur
when it is possible to write 2 or more valid
Lewis structures for the same molecule or ion.
46- Experimental data indicate that the 2 bonds in
ozone are the same __length_____, BUT double
bonds are shorter than single bonds! So the
explanation for this is that the actual bonds are
__combination_______ of those in the 2 resonance
structures. The extra electron pair in ozone is
delocalized over the two bonding regions. So
each bond spends about half the time being single
and half being double.
47Expanded Octets
- Some molecules do not follow the octet rule. The
central atom has more than 8 electrons, called an
expanded octet. - In these cases, the standard method for
determining the Lewis structure will fail.
48How to figure out a Lewis structure for an
expanded octet
- 1. Calculate the available (A) number of
electrons. - 2. Give the surrounding atoms an octet, and
assume only single bonds to the central atom - 3. Place any remaining electrons as lone pairs
on the central atom so that A electrons are
included. - not always the case, but you should still get
the right answer
49Examples
F F P F F F
N A S
5(5x7) 40
50VSEPR Theory
- Valence Shell Electron Pair Repulsion Theory
- Electron pairs orient themselves in order to
minimize repulsive forces.
51A. VSEPR Theory
- Types of e- Pairs
- Bonding pairs - form bonds
- Lone pairs - nonbonding e-
52A. VSEPR Theory
- Lone pairs reduce the bond angle between atoms.
53B. Determining Molecular Shape
- Draw the Lewis Diagram.
- Tally up e- pairs on central atom.
- double/triple bonds ONE pair
- Shape is determined by the of bonding pairs and
lone pairs.
54Common Molecular Shapes
2 bond 0 lone
LINEAR 180
55Common Molecular Shapes
2 bond 1 lone
BENT lt120
56Common Molecular Shapes
2 bond 2 lone
BENT 109.5
57Common Molecular Shapes
2 bond 3 lone
Linear
58Common Molecular Shapes
3 bond 0 lone
TRIGONAL PLANAR 120
59Common Molecular Shapes
3 bond 1 lone
TRIGONAL PYRAMIDAL 107
60Common Molecular Shapes
3 bond 2 lone
T-shaped
61Common Molecular Shapes
4 bond 0 lone
TETRAHEDRAL 109.5
62Common Molecular Shapes
4 bond 1 lone
See-Saw
63Common Molecular Shapes
4 bond 2 lone
Square planar
64Common Molecular Shapes
5 bond 0 lone
Trigonal bipyramidal
65Common Molecular Shapes
5 bond 1 lone
Square pyramidal
66Common Molecular Shapes
6 bond 0 lone
Octahedral
67 Examples
3 bond 1 lone
PYRAMIDAL 107
68Examples
2 total 2 bond 0 lone
LINEAR 180
69Examples
- Use VSEPR Theory to predict the shape of
- CO2 b) ClO3- c) NO3-
- d) SCl2 e) PCl3
70UNDERSTANDING MOLECULAR GEOMETRIES
PREDICTED BY VSEPR
H
H
C
H
H
FROM THE PERSPECTIVE OF
HYBRIDIZATION OF ATOMIC ORBITALS
71Electron Both wave and particle
72Electron Both wave and particle
73s, p, d Orbitals
p Orbitals
Px
Px
Pz
Pz
Py
Py
74Hybrid Orbitals
- __Hybrid_______ orbitals are formed when several
atomic orbitals blend to form the same total
number of equivalent orbitals. This process is
known as hybridization. - Hybridization explains the tetrahedral structure
of methane. The s and px, py, and pz orbitals of
the carbon atom blend to form four identical
hybrid orbitals
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76Number of regions of high electron density Orbitals that blend to make the hybrids Hybridization
2 s p sp
3 s p p sp2
4 s p p p sp3
5 s p p p d sp3d
6 s p p p d d sp3d2
77Now see if you can use what you have learned to
predict the hybridization of some other
compounds. Lets start with water.
Lewis Diagram
of electron clouds
four
sp3 hybridization
Hybridization
78Now try carbon dioxide
Lewis Diagram
of electron clouds
Two
sp hybridization
Hybridization
Remember double or triple bonds count as 1
electron cloud in VSEPR theory.
79Now try the sulfate ion (SO4-2)
Lewis Diagram
of electron clouds
Four
sp3 hybridization
Hybridization
80Now try the carbonate ion (CO3-2)
Lewis Diagram
of electron clouds
Three
sp2 hybridization
Hybridization
Remember double or triple bonds count as 1
electron cloud in VSEPR theory.
81Now try the sulfur hexaflouride (SF6)
Lewis Diagram
of electron clouds
Six
sp3d2 hybridization
Hybridization
82Sigma and Pi Bonds
- When two atomic orbitals combine along the axis
connecting the 2 atomic nuclei (end-to-end
overlap), a __sigma________ bond is formed. One
sigma bond is formed in all bonds (single,
double, triple). - When two atomic orbitals combine with a
side-to-side overlap, a ___pi________ bond is
formed. A pi bond occurs in double and triple
bonds only.
83Ethylene (Ethene) , C2H4
84Summary
- single bond
- composed of one _sigma___ bond
- double bond
- composed of one __sigma___ and one ____pi___ bond
- triple bond
- composed of one _sigma_____ and two ___pi____
bonds
85Polarity of Bonds
- The bonding pairs of electrons in covalent bonds
are located between the ____elements_______ of
the atoms sharing the electrons. - When the pull of each nucleus for the electrons
is equally strong, the electrons are shared
equally and are located on average halfway
between the two nuclei. Recall that this type of
bond is a ___non polar__________ covalent bond. - When the pull of one nucleus for the electrons is
stronger than the other, the electrons spend more
time closer to the more electronegative nucleus.
Recall that this type of bond is a
__polar________ covalent bond.
86- The more electronegative atom acquires a partial
___negative___________ charge. The less
electronegative atom therefore acquires a partial
__positive___________ charge. - These partial charges are indicated by the
following symbols - d - for the more electronegative element.
- d for the less electronegative element.
87Ex H2O
d
d
d -
The water molecule contains two dipole moments,
sometimes just called dipoles. The dipole
moments can be drawn in with arrow notation
88Polarity of Molecules
- Note that just because a molecule contains polar
bonds, it may or may not be classifed as a polar
molecule. - The polarity of a molecule will depend on
- (1) existence of polar bonds
- (if none of the bonds are polar, the molecule is
nonpolar) - (2) shape of the molecule
- (3) orientation of the polar bonds
89- Note to help determine if the molecule will be
polar, look for lines of symmetry. - If the molecule is symmetrical
- then the dipoles will cancel and the molecule is
nonpolar. - If the dipoles do not cancel
- the molecule is polar.
90Examples