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Transcript of Chapter Eight BONDING: GENERAL CONCEPTS. Chapter 8 | Slide 2 Copyright © Houghton Mifflin Company....
![Page 1: Chapter Eight BONDING: GENERAL CONCEPTS. Chapter 8 | Slide 2 Copyright © Houghton Mifflin Company. All rights reserved. Questions to Consider What is.](https://reader036.fdocument.pub/reader036/viewer/2022081519/56649d815503460f94a66ca6/html5/thumbnails/1.jpg)
Chapter Eight
BONDING: GENERAL CONCEPTS
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 2
Questions to Consider
• What is meant by the term “chemical bond”?
• Why do atoms bond with each other to form compounds?
• How do atoms bond with each other to form compounds?
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 3
A Chemical Bond
• No simple way to define this.
• Forces that hold groups of atoms together and make them function as a unit.
• A bond will form if the energy of the aggregate is lower than that of the separated atoms.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 4
The Interaction of Two Hydrogen Atoms
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 5
The Interaction of Two Hydrogen Atoms
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 6
Key Ideas in Bonding
• Ionic Bonding – electrons are transferred
• Covalent Bonding – electrons are shared equally
• What about intermediate cases?
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 7
Polar Covalent Bond
• Unequal sharing of electrons between atoms in a molecule.
• Results in a charge separation in the bond (partial positive and partial negative charge).
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 8
The Effect of an Electric Field on Hydrogen Fluoride Molecules
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 9
Polar Molecules
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 10
Concept Check
• What is meant by the term “chemical bond?”• Why do atoms bond with each other to form
molecules?• How do atoms bond with each other to form
molecules?
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 11
Electronegativity
• The ability of an atom in a molecule to attract shared electrons to itself.
• On the periodic table, electronegativity generally increases across a period and decreases down a group.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 12
The Pauling Electronegativity Values
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 13
Concept Check
• If lithium and fluorine react, which has more attraction for an electron? Why?
• In a bond between fluorine and iodine, which has more attraction for an electron? Why?
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 14
Concept Check
• What is the general trend for electronegativity across rows and down columns on the periodic table?
• Explain the trend.
8.2
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 15
The Relationship Between Electronegativity and Bond Type
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 16
Exercise
Arrange the following bonds from most to least polar:
a) N-F O-F C-F
b) C-F N-O Si-F
c) Cl-Cl B-Cl S-Cl
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 17
Concept Check
Which of the following bonds would be the least polar yet still be considered polar covalent?
Mg-O C-O O-O Si-O N-O
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 18
Concept Check
Which of the following bonds would be the most polar without being considered ionic?
Mg-O C-O O-O Si-O N-O
8.2
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 19
Dipole Moment
• Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
• Use an arrow to represent a dipole moment.
– Point to the negative charge center with the tail of the arrow indicating the positive center of charge.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 20
Dipole Moment
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 21
No Net Dipole Moment (Dipoles Cancel)
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 22
Stable Compounds
• Atoms in stable compounds usually have a noble gas electron configuration.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 23
Isoelectronic Series
• A series of ions/atoms containing the same number of electrons.
O2-, F-, Ne, Na+, Mg2+, and Al3+
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 24
Concept Check
Choose an alkali metal, an alkaline metal, a noble gas, and a halogen so that they constitute an isoelectronic series when the metals and halogen are written as their most stable ions.
– What is the electron configuration for each species?– Determine the number of electrons for each species.– Determine the number of protons for each species.– Rank the species according to increasing radius.– Rank the species according to increasing ionization energy.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 25
Ionic Radii
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 26
We can “read” from the periodic table:
• Trends for:– Atomic size, ion radius, ionization energy,
electronegativity
• Electron configurations
• Formula prediction for ionic compounds
• Covalent bond polarity ranking
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 27
Lattice Energy
• The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.
Lattice energy = k(Q1Q2/r)
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 28
Born-Haber Cycle for NaCl
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 29
Formation of an Ionic Solid
1.Sublimation of the solid metal• M(s) M(g) [endothermic]
2.Ionization of the metal atoms• M(g) M+(g) + e [endothermic]
3.Dissociation of the nonmetal• 1/2X2(g) X(g) [endothermic]
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 30
Formation of an Ionic Solid (continued)
4. Formation of X ions in the gas phase:• X(g) + e X(g) [exothermic]
5. Formation of the solid MX• M+(g) + X(g) MX(s)
[quite exothermic]
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 31
Comparing Energy Changes
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 32
Partial Ionic Character of Covalent Bonds
• None of the bonds reaches 100% ionic character even with compounds that have a large electronegativity difference (when tested in the gas phase).
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 33
The relationship between the ionic character of a covalent bond and the electronegativity
difference of the bonded atoms
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 34
Ionic Compound
• Any compound that conducts an electric current when melted will be classified as ionic.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 35
Models
• Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 36
Fundamental Properties of Models
1. A model does not equal reality.
2. Models are oversimplifications, and are therefore often wrong.
3. Models become more complicated and are modified as they age.
4. We must understand the underlying assumptions in a model so that we don’t misuse it.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 37
Bond Energies
• To break bonds, energy must be added to the system (endothermic).
• To form bonds, energy is released (exothermic).
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 38
Bond Energies
ΔH = ΣD(bonds broken) – ΣD(bonds formed)
D represents the bond energy per mole of bonds (always has a positive sign).
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 39
Localized Electron Model
• A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 40
Localized Electron Model
• Electron pairs are assumed to be localized on a particular atom or in the space between two atoms: Lone pairs – pairs of electrons localized on
an atom Bonding pairs – pairs of electrons found in
the space between the atoms
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 41
Localized Electron Model
1. Description of valence electron arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to share electrons or hold lone pairs.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 42
Lewis Structure
• Shows how valence electrons are arranged among atoms in a molecule.
• Reflects central idea that stability of a compound relates to noble gas electron configuration.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 43
Writing Lewis Structures
1. Sum the valence electrons.
2. Place bonding electrons between pairs of atoms.
3. Atoms usually have noble gas configurations. Arrange remaining electrons to satisfy the octet rule (or duet rule for hydrogen).
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 44
Concept Check
Draw a Lewis structure for each of the following molecules:
H2 F2
HF
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 45
Concept Check
Draw a Lewis structure for each of the following molecules:
H2O
NH3
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 46
Exceptions
• When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, place the extra electrons on the central atom.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 47
Resonance
• More than one valid Lewis structure can be written for a particular molecule.
• Actual structure is an average of the resonance structures.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 48
Concept Check
Draw a Lewis structure for each of the following molecules:
CO CO2
CH3OH BF3
PCl5 NO3-
SF6
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 49
Formal Charge
• Nonequivalent Lewis structures.
• Atoms in molecules try to achieve formal charges as close to zero as possible.
• Any negative formal charges are expected to reside on the most electronegative atoms.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 50
Formal Charge
• Formal charge = # valence e– on free atom – # valence e– assigned to the atom in the molecule.
• Assume:– Lone pair electrons belong entirely to the
atom in question– Shared electrons are divided equally
between the two sharing atoms
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 51
VSEPR Model
• VSEPR: Valence Shell Electron-Pair Repulsion.
• The structure around a given atom is determined principally by minimizing electron pair repulsions.
8.13
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 52
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put electron pairs as far apart as possible.
3. Determine positions of atoms from the way electron pairs are shared.
4. Determine the name of molecular structure from positions of the atoms.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 53
VSEPR
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 54
VSEPR: Two Electron Pairs
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 55
VSEPR: Three Electron Pairs
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 56
VSEPR: Four Electron Pairs
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 57
VSEPR: Iodine Pentafluoride
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 58
Concept Check
Determine the shape for each of the following molecules, and include bond angles:
HCN
PH3
SF4
O3
KrF4
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 59
Concept Check
True or false:
A molecule that has polar bonds will always be polar.
-If true, explain why.
-If false, provide a counter-example.
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Copyright © Houghton Mifflin Company. All rights reserved.Chapter 8 | Slide 60
Concept Check
True or false:
Lone pairs make a molecule polar.
-If true, explain why.
-If false, provide a counter-example.
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