The Shapes of Molecules

Июль 22, 2022

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The Shapes of Molecules

Содержание

2. The Shapes of Molecules 10.1 Depicting Molecules and Ions with Lewis Structures 10.2 Using Lewis Structures
3. On the Value of Lewis Structures A Lewis structure is a two-dimensional (2D) representation of a
4. Figure 10.1 Steps to convert a molecular formula into a Lewis structure Molecular formula Atom placement
5. Molecular formula Atom placement Sum of valence e- Remaining valence e- Lewis structure For NF3 N
6. SAMPLE PROBLEM 10.1 Writing Lewis Structures for Molecules with One Central Atom SOLUTION: Step 1: Carbon
7. SAMPLE PROBLEM 10.2 Writing Lewis Structures for Molecules with More than One Central Atom SOLUTION: Hydrogen
8. Lewis Structures for Molecules with Multiple Bonds After applying Steps 1-4, there may not be enough
9. SAMPLE PROBLEM 10.3 Writing Lewis Structures for Molecules with Multiple Bonds PLAN: SOLUTION: PROBLEM: Write Lewis
10. Resonance: Delocalized Electron-Pair Bonding Resonance structures have the same relative placement of atoms but different locations
11. Resonance structures are not real bonding depictions. The actual molecule is a resonance hybrid, an average
12. SAMPLE PROBLEM 10.4 Writing Resonance Structures PLAN: SOLUTION: PROBLEM: Write resonance structures for the nitrate anion,
13. When two or more unsymmetrical resonance forms exist: How do you determine which form exerts the
14. Formal Charge: Selecting the Best Resonance Structure An atom “owns” all of its non-bonding electrons and
15. Resonance (continued) Smaller formal charges (either positive or negative) are preferable to larger formal charges. Avoid
16. EXAMPLE: NCO- has three possible resonance forms. Resonance (continued) Formal charges: -2 0 +1 -1 0
17. Lewis Structures for Exceptions to the Octet Rule (a) Electron-Deficient Molecules: gaseous molecules containing either Be
18. SAMPLE PROBLEM 10.5 Writing Lewis Structures for Exceptions to the Octet Rule PLAN: SOLUTION: Draw the
19. Heats of Reactions from Lewis Structures and Bond Energies (1) Break all bonds found in the
20. Figure 10.2 Using bond energies to calculate ∆H orxn ΔHorxn = ΔHoreactant bonds broken + ΔHoproduct
21. Figure 10.3 Using bond energies to calculate ΔH orxn of methane combustion Enthalpy,H BOND BREAKAGE 4
22. SAMPLE PROBLEM 10.6 Calculating Enthalpy Changes from Bond Energies SOLUTION: bonds broken bonds formed
23. SAMPLE PROBLEM 10.6 (continued) bonds broken bonds formed 4 C-H = 4 mol (413 kJ/mol) =
24. Valence-shell Electron-Pair Repulsion (VSEPR) Theory A method to predict the shapes of molecules from their electronic
25. Figure 10.4 A balloon analogy for the mutual repulsion of electron groups
26. Figure 10.5 Electron-group repulsions and the five basic molecular shapes Ideal bond angles are shown for
27. Figure 10.6 The single molecular shape of the linear electron-group arrangement Examples: CS2, HCN, BeF2
28. Figure 10.7 The two molecular shapes of the trigonal planar electron-group arrangement Examples: SO3, BF3, NO3-,
29. Factors Affecting Actual Bond Angles Observed bond angles are consistent with theoretical angles when (a) the
30. Figure 10.8 The three molecular shapes of the tetrahedral electron-group arrangement Examples: CH4, SiCl4, SO42-, ClO4-
31. Figure 10.9 Lewis structures and molecular shapes
32. Figure 10.10 The four molecular shapes of the trigonal bipyramidal electron-group arrangement Examples: SF4 XeO2F2 IF4+
33. General trend for electron-pair repulsions for similar molecules with a given electron-group arrangement: Lone pair -
34. Figure 10.11 The three molecular shapes of the octahedral electron-group arrangement Examples: SF6 IOF5 Examples: BrF5
35. Figure 10.12 The steps in determining a molecular shape Molecular formula Lewis structure Electron-group arrangement Bond
36. SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups The shape is
37. SAMPLE PROBLEM 10.7 (continued) (b) For COCl2, C has the lowest EN and will be the
38. SAMPLE PROBLEM 10.8 Predicting Molecular Shapes with Five or Six Electron Groups (b) BrF5 - 42
39. Molecular Shapes With More Than One Central Atom Combinations of the molecular shapes observed when a
40. Figure 10.13 The tetrahedral centers of ethane
41. Figure 10.13 The tetrahedral centers of ethanol
42. SAMPLE PROBLEM 10.9 Predicting Molecular Shapes with More Than One Central Atom SOLUTION:
43. Molecular Polarity Both shape and bond polarity determine molecular polarity. Dipole moment (μ) = the product
44. Figure 10.14 The orientation of polar molecules in an electric field
45. SAMPLE PROBLEM 10.10 Predicting the Polarity of Molecules (a) ammonia, NH3 (b) boron trifluoride, BF3 (c)
46. SAMPLE PROBLEM 10.10 (continued) (b) BF3 has 24 valence electrons and all electrons around the B
47. The Complementary Shapes of an Enzyme and Its Substrate
48. Biological Receptors: Olfactory Biochemistry
49. Shapes of Some Olfactory Receptor Sites Three of the proposed seven olfactory receptors having different shapes
51. Скачать презентацию

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The Shapes of Molecules
10.1 Depicting Molecules and Ions with Lewis Structures
10.2

Using Lewis Structures and Bond Energies to Calculate
Heats of Reaction

10.3 Valence-Shell Electron-Pair Repulsion (VSEPR) Theory and
Molecular Shape

10.4 Molecular Shape and Molecular Polarity

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On the Value of Lewis Structures
A Lewis structure is a two-dimensional

(2D) representation
of a molecule.

Lewis structures are used in conjunction with valence shell
electron-pair repulsion (VSEPR) theory to predict the three-
dimensional (3D) shapes of molecules.

We first consider Lewis structures for molecules with single
bonds (bond order = 1).

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Figure 10.1
Steps to convert a molecular formula into a Lewis structure
Molecular

formula

Atom placement

Sum of valence e-

Remaining valence e-

Lewis structure

Place the atom with the lowest EN in the center

Add A-group numbers

Draw single bonds and subtract 2e- for each bond

Give each atom 8e-
(2e- for H)

Step 1

Step 2

Step 3

Step 4

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Molecular formula
Atom placement
Sum of valence e-
Remaining valence e-
Lewis structure
For NF3
N
F
F
F
N 5

valence e-

F 7 e-

X 3 = 21 valence e-

Total of 26 valence e-

N is less electronegative than F; N is the central atom

Three single bonds = 6 e-

20 remaining valence e-; 6 e-
on each F, 2 e- on N (10 lone-
pairs of electrons)

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SAMPLE PROBLEM 10.1
Writing Lewis Structures for Molecules with One Central Atom
SOLUTION:
Step

1: Carbon has the lowest EN and is the central atom.
The four remaining atoms are placed around it.

Steps 2-4:
C has 4 valence e-, Cl and F each have 7. The
sum is 4 + 4(7) = 32 valence e-.

Make bonds and fill in the remaining valence
electrons, placing 8e- around each atom.

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SAMPLE PROBLEM 10.2
Writing Lewis Structures for Molecules with More than One

Central Atom

SOLUTION:

Hydrogen can have only one bond. Thus, C and O must be next to each other, with H filling in the bonds.
There are 4(1) + 1(4) + 1(6) = 14 valence electrons.
C has 4 bonds and O has 2. O has two pairs of unshared e-.

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Lewis Structures for Molecules with Multiple Bonds
After applying Steps 1-4, there

may not be enough electrons for
the central atom (or one of the central atoms) to attain an octet.
This situation suggests that a multiple bond (bond order of 2 or 3)
is present in the molecule.

STEP 5: If, after Step 4, a central atom still does not have an octet,
make a multiple bond by changing a lone-pair from one of the
surrounding atoms into a bonding pair to the central atom.

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SAMPLE PROBLEM 10.3
Writing Lewis Structures for Molecules with Multiple Bonds
PLAN:
SOLUTION:
PROBLEM:
Write Lewis

structures for the following:
(a) Ethylene (C2H4), an important reactant in the manufacture of polymers
(b) Nitrogen (N2), the most abundant atmospheric gas

For molecules with multiple bonds, Step 5 follows the other steps in Lewis structure construction. If a central atom does not have 8 e- (an octet), then electrons can be moved to form a multiple bond.

(a) There are 2(4) + 4(1) = 12 valence electrons. H can have only one bond per atom.

(b) N2 has 2(5) = 10 valence electrons. Therefore, a triple bond is required to make the octet around each N.

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Resonance: Delocalized Electron-Pair Bonding
Resonance structures have the same relative placement of

atoms but different locations of bonding and non-bonding electron pairs.

O3 can be drawn in two ways:

Neither structure is actually correct but can be redrawn to represent a structure that is a hybrid of the two - a resonance structure.

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Resonance structures are not real bonding depictions.
The actual molecule is

a resonance hybrid, an average
of the resonance forms.

For O3, two of the electron pairs (one bonding, one non-
bonding) are delocalized (i.e., their density is spread over
the entire molecule). This effect yields two identical O-O
bonds, each consisting of a single bond (localized electron
pair) and a partial double bond (from one of the delocalized
electron pairs). Resonance effects lead to fractional bond
orders.

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SAMPLE PROBLEM 10.4
Writing Resonance Structures
PLAN:
SOLUTION:
PROBLEM:
Write resonance structures for the nitrate anion,

NO3-.

After Steps 1-4, apply Step 5. Then determine if other structures can be drawn in which the electrons can be delocalized over more than two atoms.

Nitrate has 1(5) + 3(6) + 1 = 24 valence electrons.

N does not have an octet; a pair of e- is used to form a double bond.

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When two or more unsymmetrical resonance forms
exist: How do you determine

which form exerts the
most influence on the resonance hybrid?

Because the resonance hybrid is an average of the resonance
forms, one form may contribute more than the others and
“weight” the average in its favor.

Calculating formal charge in resonance forms

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Formal Charge: Selecting the Best Resonance Structure
An atom “owns” all of

its non-bonding electrons and half of its bonding electrons.

Formal charge of atom =
# valence e- - (# unshared electrons + 1/2 # shared electrons)

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Resonance (continued)
Smaller formal charges (either positive or negative) are preferable to

larger formal charges.

Avoid like charges (+ + or - - ) on adjacent atoms.

A more negative formal charge should reside on an atom with a larger EN value.

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EXAMPLE: NCO- has three possible resonance forms.
Resonance (continued)
Formal charges:
-2
0
+1
-1
0
0
0
0
-1
Forms B and

C have negative formal charges on N and O. These forms are more important than Form A.

Form C has a negative charge on O which is more electronegative than N. Therefore, Form C contributes the most to the resonance hybrid.

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Lewis Structures for Exceptions to the Octet Rule
(a) Electron-Deficient Molecules: gaseous

molecules containing
either Be or B as the central atom; have fewer than 8 electrons
around the Be or B (4 e- around Be and 6 e- around B) (BF3).

(b) Odd-Electron Molecules: have an odd number of valence
electrons; examples include free radicals, which contain a lone
(unpaired) electron and are paramagnetic (use formal charges to
locate the lone electron) (NO2).

(c) Expanded Valence Shells: for molecules that have more than
8 electrons around the central atom; use empty outer d orbitals;
occurs only with a central atom from Period 3 or higher (SF6, PCl5).

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SAMPLE PROBLEM 10.5
Writing Lewis Structures for Exceptions to the Octet Rule
PLAN:
SOLUTION:
Draw

the Lewis structures for the molecule and determine if there is an element that is an exception to the “octet rule”. Note that (a) contains P which is a Period-3 element and can have an expanded valence shell.

(a) H3PO4 has two resonance forms, and formal charges indicate the more important form.

-1

lower formal charges

more stable

(b) BFCl2 has only one Lewis structure.

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Heats of Reactions from Lewis Structures and Bond Energies
(1) Break all

bonds found in the reactants to give free atoms

(2) Reform new bonds to the free atoms to give the products

Procedure

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Figure 10.2
Using bond energies to calculate ∆H orxn
ΔHorxn = ΔHoreactant bonds

broken + ΔHoproduct bonds formed

ΔHo1 = + sum of BE

ΔHo2 = - sum of BE

ΔHorxn

Enthalpy, ∆H

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Figure 10.3
Using bond energies to calculate ΔH orxn of methane combustion
Enthalpy,H
BOND

BREAKAGE

4 BE(C-H) = +1652 kJ

2 BE(O2) = + 996 kJ

ΔHo (bond-breaking) = +2648 kJ

BOND FORMATION

4 [-BE(O-H)] = -1868 kJ

ΔHo (bond forming) = -3466 kJ

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SAMPLE PROBLEM 10.6
Calculating Enthalpy Changes from Bond Energies
SOLUTION:
bonds broken
bonds formed

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SAMPLE PROBLEM 10.6
(continued)
bonds broken
bonds formed
4 C-H = 4 mol (413 kJ/mol)

= 1652 kJ

3 Cl-Cl = 3 mol (243 kJ/mol) = 729 kJ

ΔHobonds broken = 2381 kJ

3 C-Cl = 3 mol (-339 kJ/mol) = -1017 kJ

1 C-H = 1 mol (-413 kJ/mol) = -413 kJ

ΔHobonds formed = -2711 kJ

3 H-Cl = 3 mol (-427 kJ/mol) = -1281 kJ

ΔHoreaction = ΔHobonds broken + ΔHobonds formed = 2381 kJ + (-2711 kJ) = - 330 kJ

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Valence-shell Electron-Pair Repulsion (VSEPR) Theory
A method to predict the shapes of

molecules from their electronic structures (Lewis structures do not depict
shape)

Basic principle: each group of valence electrons around a central
atom is located as far away as possible from the others in order to
minimize repulsions

Both bonding and non-bonding valence electrons around
the central atom are considered.

AXmEn symbolism: A = central atom, X = surrounding atoms,
E = non-bonding electrons (usually a lone pair)

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Figure 10.4
A balloon analogy for the mutual repulsion of electron groups

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Figure 10.5
Electron-group repulsions and the five basic molecular shapes
Ideal bond angles

are shown for each shape.

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Figure 10.6
The single molecular shape of the linear
electron-group arrangement
Examples:
CS2, HCN,

BeF2

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Figure 10.7
The two molecular shapes of the trigonal planar
electron-group arrangement
Examples:
SO3,

BF3, NO3-, CO32-

Examples:
SO2, O3, PbCl2, SnBr2

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Factors Affecting Actual Bond Angles
Observed bond angles are consistent with theoretical

angles when (a) the atoms attached to the central atom are the same and (b) when all electrons are bonding electrons of the same order.

ideal

120o

larger EN

greater electron density

Lone pairs (unshared electron pairs) repel bonding pairs more strongly than bonding pairs repel each other.

95o

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Figure 10.8
The three molecular shapes of the tetrahedral
electron-group arrangement
Examples:
CH4, SiCl4,

SO42-, ClO4-

Examples:
NH3
PF3
ClO3
H3O+

Examples:
H2O
OF2
SCl2

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Figure 10.9
Lewis structures and molecular shapes

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Figure 10.10
The four molecular shapes of the trigonal bipyramidal electron-group arrangement
Examples:
SF4
XeO2F2
IF4+
IO2F2-
Examples:
ClF3
BrF3
Examples:
XeF2
I3-
IF2-
Examples:
PF5
AsF5
SOF4

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General trend for electron-pair repulsions for similar molecules
with a given electron-group

arrangement:

Lone pair - lone pair > lone pair - bonding pair > bonding
pair - bonding pair

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Figure 10.11
The three molecular shapes of the octahedral
electron-group arrangement
Examples:
SF6
IOF5
Examples:
BrF5
TeF5-
XeOF4
Examples:
XeF4
ICl4-

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Figure 10.12
The steps in determining a molecular shape
Molecular formula
Lewis structure
Electron-group arrangement
Bond

angles

Molecular shape (AXmEn)

Count all e- groups around the central atom A

Note lone pairs and
double bonds

Count bonding and non-bonding e- groups separately.

Step 1

Step 2

Step 3

Step 4

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SAMPLE PROBLEM 10.7
Predicting Molecular Shapes with Two, Three, or Four Electron

Groups

The shape is based on the tetrahedral arrangement.

The F-P-F bond angles should be < 109.5o due to the repulsion of the non-bonding electron pair.

The final shape is trigonal pyramidal.

< 109.5o

The type of shape is:
AX3E

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SAMPLE PROBLEM 10.7
(continued)
(b) For COCl2, C has the lowest EN and

will be the center atom.

There are 24 valence e-, with 3 atoms attached to the center atom.

C does not have an octet; a pair of non-bonding electrons will move in from the O to produce a double bond.

The shape for an atom with three atom attachments and no non-bonding pairs on the central atom is trigonal planar.

The Cl-C-Cl bond angle will be less than 120o due to the electron density of the C=O.

Type AX3

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SAMPLE PROBLEM 10.8
Predicting Molecular Shapes with Five or Six Electron Groups
(b)

BrF5 - 42 valence e-; 5 bonding pairs and 1 non-bonding pair on the central atom. Shape is AX5E, square pyramidal.

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Molecular Shapes With More Than One Central Atom
Combinations of the molecular

shapes observed
when a single central atom is present

Examples: CH3-CH3 (ethane) and CH3CH2OH (ethanol)

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Figure 10.13
The tetrahedral centers of ethane

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Figure 10.13
The tetrahedral centers of ethanol

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SAMPLE PROBLEM 10.9
Predicting Molecular Shapes with More Than One Central Atom

SOLUTION:

Слайд 43

Molecular Polarity
Both shape and bond polarity determine molecular polarity.
Dipole moment (μ)

= the product of the partial charges caused by
polar bonds and the distances between them; debye (D) units,
where 1 D = 3.34 x 10-30 coulomb . meter

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Figure 10.14
The orientation of polar molecules in an electric field

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SAMPLE PROBLEM 10.10
Predicting the Polarity of Molecules
(a) ammonia, NH3
(b) boron trifluoride,

BF3

ENN = 3.0

ENH = 2.1

bond dipoles

molecular dipole

The bond dipoles reinforce each other, so the overall molecule is polar.

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SAMPLE PROBLEM 10.10
(continued)
(b) BF3 has 24 valence electrons and all electrons

around the B will be involved in bonds. The shape is AX3 (trigonal planar).

F (EN 4.0) is more electronegative than B (EN 2.0) and all of the bond dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is non-polar.

120o

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The Complementary Shapes of an Enzyme and Its Substrate

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Biological Receptors: Olfactory Biochemistry

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Shapes of Some Olfactory Receptor Sites
Three of the proposed seven olfactory

receptors
having different shapes

The Shapes of Molecules

Содержание

The Shapes of Molecules10.1 Depicting Molecules and Ions with Lewis Structures10.2

On the Value of Lewis StructuresA Lewis structure is a two-dimensional

Figure 10.1Steps to convert a molecular formula into a Lewis structureMolecular

Molecular formulaAtom placementSum of valence e-Remaining valence e-Lewis structureFor NF3NFFFN 5

SAMPLE PROBLEM 10.1Writing Lewis Structures for Molecules with One Central AtomSOLUTION:Step

SAMPLE PROBLEM 10.2Writing Lewis Structures for Molecules with More than One

Lewis Structures for Molecules with Multiple BondsAfter applying Steps 1-4, there

SAMPLE PROBLEM 10.3Writing Lewis Structures for Molecules with Multiple BondsPLAN:SOLUTION:PROBLEM:Write Lewis

Resonance: Delocalized Electron-Pair BondingResonance structures have the same relative placement of

Resonance structures are not real bonding depictions. The actual molecule is

SAMPLE PROBLEM 10.4Writing Resonance StructuresPLAN:SOLUTION:PROBLEM:Write resonance structures for the nitrate anion,

When two or more unsymmetrical resonance formsexist: How do you determine

Formal Charge: Selecting the Best Resonance StructureAn atom “owns” all of

Resonance (continued)Smaller formal charges (either positive or negative) are preferable to

EXAMPLE: NCO- has three possible resonance forms.Resonance (continued)Formal charges:-20+1-10000-1Forms B and

Lewis Structures for Exceptions to the Octet Rule(a) Electron-Deficient Molecules: gaseous

SAMPLE PROBLEM 10.5Writing Lewis Structures for Exceptions to the Octet RulePLAN:SOLUTION:Draw

Heats of Reactions from Lewis Structures and Bond Energies(1) Break all

Figure 10.2Using bond energies to calculate ∆H orxnΔHorxn = ΔHoreactant bonds

Figure 10.3Using bond energies to calculate ΔH orxn of methane combustionEnthalpy,HBOND

SAMPLE PROBLEM 10.6Calculating Enthalpy Changes from Bond EnergiesSOLUTION:bonds brokenbonds formed

SAMPLE PROBLEM 10.6(continued)bonds brokenbonds formed4 C-H = 4 mol (413 kJ/mol)

Valence-shell Electron-Pair Repulsion (VSEPR) TheoryA method to predict the shapes of

Figure 10.4A balloon analogy for the mutual repulsion of electron groups

Figure 10.5Electron-group repulsions and the five basic molecular shapesIdeal bond angles

Figure 10.6The single molecular shape of the linear electron-group arrangementExamples:CS2, HCN,

Figure 10.7The two molecular shapes of the trigonal planar electron-group arrangementExamples:SO3,

Factors Affecting Actual Bond AnglesObserved bond angles are consistent with theoretical

Figure 10.8The three molecular shapes of the tetrahedral electron-group arrangementExamples:CH4, SiCl4,

Figure 10.9Lewis structures and molecular shapes

Figure 10.10The four molecular shapes of the trigonal bipyramidal electron-group arrangementExamples:SF4XeO2F2IF4+IO2F2-Examples:ClF3BrF3Examples:XeF2I3-IF2-Examples:PF5AsF5SOF4

General trend for electron-pair repulsions for similar moleculeswith a given electron-group

Figure 10.11The three molecular shapes of the octahedral electron-group arrangementExamples:SF6IOF5Examples:BrF5TeF5-XeOF4Examples:XeF4ICl4-

Figure 10.12The steps in determining a molecular shapeMolecular formulaLewis structureElectron-group arrangementBond

SAMPLE PROBLEM 10.7Predicting Molecular Shapes with Two, Three, or Four Electron

SAMPLE PROBLEM 10.7(continued)(b) For COCl2, C has the lowest EN and

SAMPLE PROBLEM 10.8Predicting Molecular Shapes with Five or Six Electron Groups(b)

Molecular Shapes With More Than One Central AtomCombinations of the molecular

Figure 10.13The tetrahedral centers of ethane

Figure 10.13The tetrahedral centers of ethanol

SAMPLE PROBLEM 10.9Predicting Molecular Shapes with More Than One Central Atom

Molecular PolarityBoth shape and bond polarity determine molecular polarity.Dipole moment (μ)

Figure 10.14The orientation of polar molecules in an electric field

SAMPLE PROBLEM 10.10Predicting the Polarity of Molecules(a) ammonia, NH3(b) boron trifluoride,

SAMPLE PROBLEM 10.10(continued)(b) BF3 has 24 valence electrons and all electrons