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Chapter 9 Chemical Bonding

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Chapter 9 Chemical Bonding * Mg ions and oxide ions are smaller than K ions and Cl ions. Also, Mg and O have 2+ and 2- charges, respectively. K and Cl only have ... – PowerPoint PPT presentation

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Title: Chapter 9 Chemical Bonding


1
Chapter 9Chemical Bonding
2
Section 9.1 Why does bonding occur in the first
place?
Bonding lowers the potential energy between
positive and negative particles (p341). What is
potential energy?
Energy changes forms P.E. ? Kinetic Energy
(K.E.)
3
Section 9.1 Why does bonding occur in the first
place?
Bonding lowers the potential energy between
positive and negative particles (p341).
4
Section 9.1 Why does bonding occur in the first
place?
Bonding lowers the potential energy between
positive and negative particles (p341).
When chemical bonds form Chemical P.E. changes
to Heat Energy Light Energy
Mechanical Energy
Electrical Energy
Heat (Thermal) Energy
Light (Radiant) Energy
Chemical Energy
5
Section 9.1 Why does bonding occur in the first
place?
Bonding lowers the potential energy between
positive and negative particles (p341).
Energy changes forms Chemical P.E. ? Heat
Light Energy
http//chemsite.lsrhs.net/chemKinetics/PotentialEn
ergy.html
6
Section 9.1 Three Type of Bonds
Ionic bonding Metal Nonmetal (Valence e-
transferred) Covalent bonding Nonmetal
Nonmetal (Valence e- shared) Metallic bonding
Metal Metal (Sea of e-)
http//chemed.chem.purdue.edu/genchem/topicreview/
bp/ch10/non.php
7
Concept Check
Review Valence Electrons e- involved in
forming compounds (Ch 8, p315)
Boron (B)
Magnesium (Mg) Hydrogen (H)
How many valence e-?
How many needed for full outer shell?
Total valance e-
8
Section 9.1 Two Bond Types With Localized
Electrons Ionic Covalent Bonding
9
Section 9.1 Two Bond Types With Localized
Electrons Ionic Covalent Bonding Why do
ionic bonds form instead of covalent bonds, and
vice versa?
Bonding Continuum
nonmetals nonmetal
metal nonmetal
Polar Covalent Bond
Nonpolar Covalent Bond
Ionic Bond
Electrons are shared unequally.
Electrons are transferred.
10
Extent of electron sharing in Covalent Bonds
e-s shared between atoms of the same
element Equal Sharing
e-s shared between atoms of different
elements Unequal Sharing
Unequal sharing occurs because one of the atoms
in a bond has a stronger attraction for the
pair of e-s than does the other atom
Why does one atom have a stronger attraction for
e-?
11
Electronegativity
Definition electronegativity (E.N) the
ability of an atom to attract the shared electrons
Increasing E.N.
Decreasing E.N.
Rule for Bond Formation The atom with the
greater E.N. pulls the shared electrons closer to
its nucleus resulting in (1) charge on high
E.N. atom (2) charge on low
E.N. atom
More later Section 9.5
12
Why do ionic bonds form instead of covalent
bonds, and vice versa?
Answer Electronegativity Differences
Example Oxygen (O) bonds with Magnesium (Mg)
MgO E.N. of O 3.5 E.N. of Mg 1.2 E.N.
difference 2.3
13
Section 9.1 The Other Bond Type With
Delocalized Electrons Metallic Bonding
Metallic Bonding - Delocalized
Covalent Bonding, Ionic Bonding
- Delocalized
A messy sea of electrons
Electrons fit neatly into shells.
14
Section 9.1 The Other Bond Type With
Delocalized Electrons Metallic Bonding
Metallic Bonding - Delocalized
A messy sea of electrons
15
Lewis Electron-Dot Symbols
Two parts
(1) Element symbol nucleus inner
electrons Ex The element lithium has an element
symbol Li
(2) Surrounding dots valence electrons (outer
most shell)
16
Review Ions Ion charged
particles that form when an atom gains or loses
one or more electrons (Ch2, p60)
Element
Ion
Ion Type
Mg
Cation
Cl
Anion
17
Review Electron Configuration and Orbital
Diagrams (Ch8, p304-317)
Example
18
Concept Check
End of Chapter Problems in-class (for
now) 9.7, 9.9, 9.13, 9.15 Write the ion for
the following elements K, Br, Sr, Ar, O For
example, the ion for Mg is Mg2.
Suggested Optional Practice Problems (for
outside of class) 9.6, 9.8, 9.10, 9.12, 9.14
(Answers in back of book or online)
19
Section 9.2 Ionic Bonding
Central idea Electrons are transferred from
metal atoms to nonmetal atoms to form
ions that come together in a solid
ionic compound.
Solid Ionic compound
Na metal Cl - nonmetal
Sodium chloride (NaCl)
Contrast with molecules formed during covalent
bonding (more later).
20
Section 9.2 Ionic Bonding
Rule The total number of e- lost by the metal
atom equals the total number gained by
the nonmetal atom.
Na
lost
gained
21
Behavior of Ionic Compounds
Why is the melting point of MgO higher than the
melting point of KCl?
22
Lattice Energy (?Hºlattice)
23
Section 9.2 Lattice Energy
Definition The enthalphy change that occurs
when 1 mol of ionic solid separates into
gaseous ions. For Review of Enthalpy Ch6,
p243 Lattice Energy denoted as ?Hºlattice
?Hºlattice cannot be measured directly, BUT
it can be calculate using the
Born-Haber cycle
24
Section 9.2 Born-Haber Cycle
Uses Hesss Law Total enthalpy of an overall
reaction is the sum of the enthalpy
changes of individual
reactions. (?Htotal ?Hrxn1 ?Hrxn2 .)
Not actual steps.
25
Section 9.2 Trends in Lattice Energy
Coulombs Law (Ch2)
26
Section 9.2 Trends in Lattice Energy
27
Behavior of Ionic Compounds
So, why is the melting point of MgO higher than
the melting point of KCl?
28
Concept Check
End of Chapter Problems in-class (for
now) 9.27, 9.30
Suggested Optional Practice Problems (for
outside of class) 9.26, 9.28 (Answers in back
of book)
29
Problem 9.30
30
Section 9.3 Covalent Bonding
e- sharing primary way that atoms interact
Nonmetal Nonmetal
Examples Water (H2O) Carbon Dioxide (CO2)
Organic Compounds
O
O
O
H
H
H
C
C
C
C
H
H
H
H
H
H
H
31
Section 9.3 Covalent Bonding Why do covalent
bonds form?
Lower P.E. More stable
32
Section 9.3 Covalent Bonding How are the
electrons distributed?
Electron Density
In order for each atom to have a full outer shell
(2 e- for H, He 8 e- for others), the electrons
arrange themselves in certain configurations
Bonding Pairs Lone Pairs
Bond Type double, single, triple
33
Section 9.3 Covalent Bonding Bond Energy
(B.E.) aka Bond Enthalpy or Bond Strength
Covalent Bond Strength depends on strength of
attraction between nuclei and
shared electrons
34
Section 9.3 Covalent Bonding Bond Energy (B.E.)
Bond formation is exothermic ?Hº always
Bond breakage is endothermic ?Hº always -
Absolute value of B.E. Each bond has its own
unique B.E. due to variations in (1) e-
density (2) charge (3) atomic radii
35
Section 9.3 Covalent Bonding Strength of Bond
different than E required to pull atoms apart
(B.E.)
Less E needed to break. Lower B.E.
Weaker Bonds Higher Energy Shallow Energy Well
Stronger Bonds Lower Energy Deeper Energy Well
More E needed to break. Higher B.E.
36
Section 9.3 Covalent Bonding Bond Energy
(B.E.) and Bond Length
Bond Length sum of the radii of the bonded
atoms (analogous to distance in Coloumbs Law)
At minimum E point.
37
Section 9.3 Covalent Bonding Bond Energy
(B.E.) and Bond Length
This relationship holds, in general, ONLY for
single bonds.
38
Section 9.3 Covalent Bonding Bond Type
(Single, Double, Triple) also matters
Same two elements, different B.E.
Nuclei more attracted to 2 shared pairs of e-
than one shared pair of e-. Higher bond order
Shorter bond length Higher Bond Energy
39
Section 9.3 Covalent Bonding Periodic Table
Trends Without Detailed Bond Lengths
The closer the atoms, the stronger the bond.
Bond Energy CF gt CCl gt CBr
40
Section 9.3 Covalent Bonding Covalent Bonds
are stronger than Ionic Bonds So why, then, do
covalent compounds have lower melting points than
ionic compounds?
Example CCl4 m.p. -23 ºC NaCl m.p.
800 ºC
41
Section 9.4 Bond Energy and Chemical Change
Where does the heat that is released come from?
http//chemsite.lsrhs.net/chemKinetics/PotentialEn
ergy.html
42
Section 9.4 Bond Energy and Chemical Change
Total energy of a chemical system K.E. P.E.
Example of a chemical system A container filled
with molecules.
http//www.landfood.ubc.ca/courses/fnh/301/water/m
otion.gif
43
Section 9.4 Bond Energy and Chemical Change
This leaves us with changes in P.E. during
chemical reactions.
P.E. contributions can from electrostatic forces
between Separate Vibrating Atoms Nucleus
Electrons in Atoms Protons Neutrons in
Nucleus Nuclei and Shared Electron Pair in Each
Bond
44
Section 9.4 Bond Energy and Chemical Change
Heat of reaction, ?Hºrxn
Exothermic reaction - ?Hºrxn
Endothermic reaction ?Hºrxn
?Hºrxn ?Hºreactant bonds broken ?Hºproduct
bonds formed
?Hºrxn ?BEreactant bonds broken ?BEproduct
bonds formed
45
Section 9.4 Bond Energy and Chemical Change
Example H2 F2 ? 2 HF
Weaker Bonds Less Stable, More Reactive H2 and
F2
Stronger Bond More Stable, Less Reactive HF
46
Section 9.4 Bond Energy and Chemical Change
Another way to looks at this reaction
H2 F2 ? 2 HF
Heat of reaction, ?Hºrxn
2 H 2 F
H2 F2
HF
?Hºrxn ?Hºreactant bonds broken ?Hºproduct
bonds formed
47
Section 9.4 Bond Energy and Chemical Change
Use bond energies to calculate ?Hºrxn (Table 9.2)
H2 F2 ? 2 HF
9.39, 9.47, 9.49
Optional Homework Problems 9.38, 9.46, 9.48,
9.50
48
Section 9.4 Bond Energy and Chemical Change
Application Energy Released From Combustion of
Fuel
49
Section 9.5 Between the Extremes
Scientific models are idealized descriptions of
reality.
Electronegativity the relative ability of a
bonded atom to attract the shared e-
50
Section 9.5 Between the Extremes
Electronegativity inversely related to atomic
size (radius) WHY?
51
Section 9.5 Between the Extremes
Nonmetals are more electronegative than metals.
52
Section 9.5 Between the Extremes
Electronegativity and Oxidation Number (O.N.)
(Review of O.N. Section 4.5)
Oxidation-reduction (redox) reactions The net
movement of electrons from one reactant to the
other. Oxidation the loss of e-
(LEO), Reduction the gain of e- (GER)
LEO the lions says GER!
Oxidizing agent becomes reduced Reducing
agent becomes oxidized
53
Oxidation Number and Electronegativity
When dead organisms (such as plankton) fall to
the bottom of the sea, their dead bodies are
eaten (respiration) by bacteria living in the
ocean sediments CH2O O2 ? CO2 H2O What
might be a problem for bacteria trying to eat
CH2O deep in sediments?
In addition to O2 SO42- and NO32- are present
in the sediments. Which might they use?
54
Section 9.5 Between the Extremes
  • Electronegativity
    and Oxidation Number (O.N.)
  • E.N. is used to determine an
    atoms O.N. in a given bond.
  • The more E.N. atom in a bond is assigned ALL the
    SHARED e- The less
  • E.N. atoms is assigned NONE
  • Example HCl Cl 8
    H 0
  • (2) O.N. valence e- - shared e-
  • Example O.N.Cl 7 8 -1 O.N.H
    1 0 1

55
Section 9.5 Between the Extremes
Polar Covalent Bonds
  • This bond type is indicated by
  • polar arrow ( ) pointing toward
    negative pole HF
  • delta symbol (?)

56
Section 9.5 Between the Extremes
Polar Covalent vs. Nonpolar Covalent
57
Section 9.5 Between the Extremes
Partial Ionic Character related directly to the
electronegativity difference (?EN)
Why? A greater ?EN results in larger partial
charges (?) and a higher partial ionic character.
Example HCl, LiCl, Cl2 Arrange these compounds
in order of least to most partial ionic character.
58
Section 9.5 Between the Extremes
Two approaches for getting a sense of a
compounds ionic character 1 Arbitrary
cutoffs used in bonding continuum.
59
Section 9.5 Between the Extremes
Two approaches for getting a sense of a
compounds ionic character 2 Calculate the
percent ionic character (increases with
?EN) Compare actual behavior of a polar
molecule in an electric field with the behavior
it would show if the e- were completely
transferred (pure ionic). 50 is dividing
line. Notice Cl2 is 0 ionic, but no molecule
has 100 ionic character (e- sharing occurs to
some extent in every bond.
60
Section 9.5 Between the Extremes
Notice, now Why metal that bond with nonmetals
form ionic bonds. Why nonmetals
that bond with other nonmetals form covalent
bonds.
61
Section 9.5 Between the Extremes
Properties of substances are indicative of their
ionic or covalent character.
62
Section 9.6 Metallic Bonding (More in Chap 12)
Electron Sea Model
In reactions with nonmetals, metals (Na) transfer
their outer e- to form ionic solids (NaCl).
What holds together bonded metals (Na)? All
metal atoms contribute their valence e-, which
are shared among all the atoms in a sample.
Metallic Bonding - Delocalized
Covalent Bonding, Ionic Bonding
- Localized
A messy sea of electrons
Electrons fit neatly into shells.
Alloys - more than one metal element involved in
a metallic sea
63
Section 9.6 Metallic Bonding (More in Chap 12)
Properties of metal substances are explained by
the electron sea model. Most metals are
solids. High m.p. attractions b/w cations and
anions need not be broken Much higher b.p.
attractions b/w cations and anions broken m.p.
depends on of valence e-
64
Problems for today 9.62, 9.64, 9.66 What
would you expect the B.E. of a HF bond to be
given that HH 432 kJ/mol FF 159
kJ/mol ?
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