Molecular Structures - Middle Tennessee State Universitymtweb.mtsu.edu/nchong/Chapter 9-CHEM1010-MSJ.pdf · Molecular Structures ... • lone pair/bond pair are intermediate in strength

1

John W. MooreConrad L. StanitskiPeter C. Jurs

Stephen C. Foster • Mississippi State University

http://academic.cengage.com/chemistry/moore

Chapter 9Molecular Structures

Molecular Structures

dimethyl ether

H – C – O – C – H

H|

|H

H|

|H

..

..

ethanol

H – C – C – O – H

H|

|H

H|

|H

..

..

C2H6O structural isomers:

Molecular shape is important!

Small structural changes cause large propertychanges.

m.p./ °C -114 -142b.p./ °C +78 -25

Physical models of 3D-structures:

ball and stick space filling

Computer versions:

Using Molecular Models

Hand-drawn molecules:

H

CH H

HIn the plane of

the screen

Going back intothe screen

Coming out ofthe screen

Using Molecular Models

The Valence Shell Electron Pair RepulsionValence Shell Electron Pair Repulsion model isused to predict shapes. Key ideas:

1. e- pairs stay as far apart as possible.

• Repulsions are minimized.

2. Molecule shape is governed by the number ofbond pairs and lone pairs present.

3. Treat multiple bonds like single bonds.

• Each is a single e- group.

4. Lone pairs occupy more volume than bonds.

Predicting Molecular Shapes: VSEPR Predicting Molecular Shapes: VSEPR

Linear Triangular planar Tetrahedral

Triangular bipyramidal Octahedral

2

Shapes that minimize repulsions:

linear triangularplanar

tetrahedral triangularbipyramidal

octahedral

Predicting Molecular Shapes: VSEPR

Bonds and lone pairs determine shape.Use the notation AXnEm

n atoms bonded tocentral atom A

m lone pairs oncentral atom A

If a molecule contains:

• bonding pairs only – these angles are correct:

• These angle change (a little) if any “X” is replaced by alone pair:

• lone pair/lone pair repulsions are largest.

• lone pair/bond pair are intermediate in strength.

• bond/bond interactions are the smallest.


Molecules may be described by their:

• electron-pair (e- pair) geometry

• molecular geometry (molecular shape)

These geometries may be different.

• Atoms can be “seen”, lone pairs are invisible.


2 e- groupsbond lonepairs pairs

2 0 AX2E0

linear

....1 1 AX1E1

linear

Linear e- pair geometry

molecular geometry


AXnEm: 2 e- group central atoms (m + n = 2)

Linear.

180.0°

180.0° “2” bonds, 0 lone pairs on C.(treat double bonds as 1 bond)Linear.

OCO

ClBeCl

Each H-C-C unit is linear.HCCH

180.0°

180.0°


AX2E0 examples:


........

3 0 AX3E0

triangular planar

2 1 AX2E1

angular (bent)

1 2 AX1E2

linear

Triangular planar e- pair geometry

molecular geometry


AXnEm: 3 e- group central atoms (m + n = 3)

3

AX3E0 examples:

Triangular planar.

Each C is AX3E0 = triangular planar.

ClBCl

Cl

CCH H

H H

120°



4 0 AX4E0

tetrahedral

....

AX1E3?All molecules with only1 bond are linear!

3 1 AX3E1

triangular pyramidal

2 2 AX2E2

angular (bent)

....


AXnEm: 4 e- group central atoms (m + n = 4)Tetrahedral e- pair

geometry

moleculargeometry

AX4E0

All angles = tetrahedral angle.

AX3E1

Lone-pair/bond > bond/bondrepulsion: H-N-H angle isreduced.

AX2E2

Two lone pairs: H-O-H angle iseven smaller.

HCH

H

H

HNH

H

OH

H


VSEPR applies to each atom in a molecule.• Alkanes: each C is tetrahedral.


Tetrahedral O

Lactic acid:

Tetrahedral C

Triangular planar C

Tetrahedral C

H

CCH

H

C

O

O

O

H

H

H

..

..

....

.. ..

Tetrahedral O


Bond pairs Lone pairs Shape

5 0 Triangular bipyramidal

4 1 Seesaw

3 2 T-shaped

2 3 Linear

6 0 Octahedral

5 1 Square pyramidal

4 2 Square planar

3 3 T-shaped

Central atoms with five or six e- pairs:

• lone pairs repel the most.• they get as far apart as possible.

Expanded Octets

4

The atoms are non-equivalent.

Green atoms are axialaxial; blue atoms are equatorialequatorial.

Expanded OctetsAXnEm: m + n = 5Triangular bipyramidal e- pair geometry.

Triangularbipyramidal

Seesaw T-shaped Linear

FPF

F

F F

FSF

F

F

FClF

F

FXeF

Expanded Octets

Expanded Octets

AXnEm: m + n = 6 Octahedral e- pair geometry:

Octahedral Square pyramid Square planar

All atoms areequivalent in

AX6E0 FSF

F

F F

F

FBrF

F

F F

ClICl

Cl

Cl

Expanded Octets

Lewis dot + VSEPR predict molecular shapes, butbut…

How do atomic orbitals (s, p…) lead to these shapes?

Valence bond theoryValence bond theory:: bonds occur when partially-occupied atomic orbitals overlap.

Orbitals Consistent with Molecular Shapes

H2 – H(1s) overlaps H(1s)

74 pm

HF – H(1s) overlaps F(2p)

109 pm

Valence Bond Theory

This works for H2 and HF, but why does…

• Be form compounds?

• Be (1s2 2s2).

• No unpaired e- to share.

• Experiments show: linear BeH2, BeCl2, …

• C form 4 bonds at tetrahedral angles?

• C (1s2 2s2 2p2).

• 2px1 2py

1 Two bonds?

• p orbitals are at 90° to each other

• Experiments show: tetrahedral CH4, CCl4, …

5

Atomic orbitals (AOs) can be hybridizedhybridized (mixed).

• Sets of identical hybridhybrid orbitals form identical bonds.

• Number of hybrids formed = number of AOs mixed.

One s orbital + one p orbital → two sp hybrids.

Orbitals Consistent with Molecular Shapes sp Hybrid OrbitalsBe compounds (BeH2, BeF2 …):

Each sp hybrid (180° apart) holds one e-.

Two equivalent covalent bonds form.

sp2 Hybrid OrbitalsB forms three sp2 hybrid orbitals:

• One s orbital mixes with two p orbitals.

• One p orbital remains unmixed.

sp2 Hybrid OrbitalsB compounds (BH3, BF3 …):

Each sp2 hybrid (120° apart) holds one e-.

Three equivalent covalent bonds form.

sp3 Hybrid OrbitalsC forms four sp3 hybrid orbitals:

• One s orbital mixes with three p orbitals.

• All p orbitals are mixed.

In C, each sp3 hybrid (109.5° apart) holds one e-.

Four equivalent covalent bonds form.

sp3 Hybrid OrbitalsN and O compounds (NH3, H2O…) have more e-:

6

sp3 Hybrid Orbitals“Octet rule” molecules have tetrahedral e- pair shape.

• sp3 hybridized (CH4, NH3, H2O, H2S, PH3, …)

Head-to-head bond = a sigma bondsigma bond (σσ bondbond).

There are:

• 4 σ bonds in CH4

• 3 σ bonds in NH3

• 2 σ bonds in H2O

H

C

HH

H

σ bond

Summary:

Mixed Hybrids (#) Remaining Geometry

s+p sp (2) p+p Linear

s+p+p sp2 (3) p Triangular planar

s+p+p+p sp3 (4) Tetrahedral

d orbitals can also form hybrids:

Mixed Hybrids (#) Remaining Geometry

s+p+p+p+d sp3d (5) d+d+d+d Triangular bipyramid

s+p+p+p+d+d sp3d2 (6) d+d+d Octahedral

Hybridization

Carbon atoms form:

• tetrahedral centers (CH4, CHF3 , C2H6…) = sp3

• triangular-planar centers (H2CO, C2H4 …) = sp2

CCH H

H H

The double bond in ethene is composed of:

• a σσ bondbond – head-to-head overlap of sp2 on each C atom.

• a ππ bondbond – sideways overlap of p AOs on the C atoms.

Hybridization in Molecules with Multiple Bonds

C (sp2) + C (sp2) overlap (σ bond):C C

H

H H

H

Unhybridized C p orbitals each contain one e-.

C C

H H

H

σ bondC C

H H

H

overlap

Sideways overlap forms oneone π bond• the lobes above and below the plane together equal 1 bond



Formaldehyde is similar: C also forms linear centers:

• C2H2 (acetylene) = sp hybridized

The triple bond is:

• one σσ bondbond

• two ππ bondsbonds

• sp hybridization leaves two unmixed p orbitalson each C.

CCH H


7

σ bond: C (sp) + C (sp) overlap:C C HH

TwoTwo π bonds• above and below overlaps are 1 bond.• front and back overlaps are a second bond.

TwoTwo p orbitals on eacheach C contain a single e-.

C C HH overlap C CH H


Molecule C-C bonding C-C rotation

ethane (CH3–CH3) σσ yes

ethene (CH2=CH2) σσ ++ ππ no

ethyne (HC≡CH) σσ ++ ππ ++ ππ no

π bonds prevent bond rotation:

Non-rotating double bonds allow cis-trans isomerismto occur.


• Most bonds are polar (e.g. C-O)

• O is δ-, C is δ+ (ENO = 3.5, ENC = 2.5)

• But many moleculesmolecules are nonpolar (e.g. CO2).

• The dipoles cancel because of CO2’s shape.• have equal size but point in opposite directions.

arrow points to δ-,the + shows δ+

O = C = Oδ-δ- 2δ+

Molecular Polarity

• Water is polar (bond dipoles do not cancel)

Dipole, μ = 1.85 D

H H

O+

Netdipole

Molecular Polarity

Molecular Polarity

Dipole momentDipole moment (μ) is a measure ofmolecule polarity:

Units: coulomb meter (Cm)

Debye (D)

Molecule μ (D)

H2 0

HF 1.78

HCl 1.07

HBr 0.79

HI 0.38

CH4 0

CH3Cl 1.92

CH2Cl2 1.60

CHCl3 1.04

CCl4 0nonpolar (μ=0)

highly polar

weakly polar

A molecule is nonpolarnonpolar if it is:

• AXnEE00 and all XX are identical.

CO2 AX2E0 linear

CH4 AX4E0 tetrahedral

CCl4 AX4E0 tetrahedral

PF5 AX5E0 triangular bipyramidal

• “divisible” into nonpolar AXnE0 shapes

PCl3F2 triangular planar (PCl3) + linear (PF2 )

XeF4 linear (XeF2) + linear (XeF2 )

Molecular Polarity

8

AXnEm molecules are polarpolar if they don’t divide intononpolar shapes, and::

Molecular Polarity

How polar? It depends on the number, type, andgeometry of the polar bonds.

• m ≠ 0:

H2O AX2E2 bent polarpolar

NH3 AX3E1 pyramidal polarpolar

• The X in AXnE0 differ:

CH2Cl2 AX4E0 tetrahedral polarpolarPF4Cl AX5E0 triangular bipyramidal polarpolar

Molecular Polarity

F F

F

C

F

CF4 is non polar

No net dipole

F F

H

C

F

CHF3 is polar

Netdipole

+

Non polarNon polarAX5E0; identical X

PCl5

PCl4F

Non polarNon polarAX5E0 and “X” differ.

BUT divisible intononpolar shapes:linear + triangularlinear + triangular

planarplanar

PF3Cl2

PolarPolarAX5E0

“X” differ

+

PolarPolarAX5E0 and “X” differ.Doesn’t divide intononpolar shapes

Molecular Polarity

PCl3F2

Molecules attract each other.

IntermolecularIntermolecular forces:

• also called noncovalentnoncovalent interactions.

• are small (compared to bonding forces).

• do not include ionic or metallic-bonding forces.

Three types:

• London forces.

• dipole-dipole attraction.

• hydrogen bonding.

Noncovalent Interactions

London ForcesAlso called dispersiondispersion forces.

• Random e- motion produces a temporary dipole inone molecule which induces a dipole in another.

• Strength (0.05 ↔ 40 kJ/mol):

Small molecule = few e- = weak attraction.

Large molecule = many e- = stronger attraction.

• Occur between all atoms and molecules.

The only force between nonpolar molecules.

Noble Gas Halogen Hydrocarbon

# of e- bp (°C) # of e- bp (°C) # of e- bp (°C)

He 2 −269 F2 18 −188 CH4 10 −161

Ne 10 −246 Cl2 34 −34 C2H6 18 −88

Ar 18 −186 Br2 70 +59 C3H8 26 −42

Kr 36 −152 I2 106 +184 C4H10 34 0

More e- = larger attraction = higher b.p.

London Forces

9

Polar molecules attract each other.

Strength: 5 ↔ 25 kJ/mol.

Dipole-Dipole Attractions

Nonpolar Molecules Polar Molecules

# of e- bp (°C) # of e- bp (°C)

SiH4 18 −112 PH3 18 −88

GeH4 36 −90 AsH3 36 −62

Br2 70 +59 ICl 70 +97

Relative importance of dipole/dipole and London ishard to predict:

Dipole-Dipole Attractions

Dipole London bp (°C)

HI small (0.38 D) large (54 e-) −36

HCl large (1.07 D) small (18 e-) −85

stickier

An especially large dipole-dipole attraction.

• 10 ↔ 40 kJ/mol.

• Occurs when H bonds directly to F, O or N.

F, O & N are small with large electronegativities.

• results in large δ+ and δ- values.

H-bonds are usually drawn as dotted lines.

Hydrogen Bonding

H on one moleculeinteracts with O onanother molecule.

Hydrogen Bonding

Water is a liquid atroom T (not a gas).

Hydrogen Bonding

Documents

Molecular Structures - Middle Tennessee State Universitymtweb.mtsu.edu/nchong/Chapter 9-CHEM1010-MSJ.pdf · Molecular Structures ... • lone pair/bond pair are intermediate in strength