Molecular Structures. Chapter 9 Molecular Structures. Using Molecular Models. Using Molecular Models. C 2 H 6 O structural isomers: .. H C C O..

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John W. Moore Conrad L. Stanitski Peter C. Jurs

Stephen C. Foster • Mississippi State University

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

Chapter 9 Molecular 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 property changes.

m.p./ °C -114 -142

b.p./ °C +78 -25

Physical models of 3D-structures:

ball and stick space filling

Computer versions:

Using Molecular Models

Hand-drawn molecules:

H C

H H

In the plane of H

the screen

Going back into the screen

Coming out of the screen

Using Molecular Models

The Valence Shell Electron Pair RepulsionValence Shell Electron Pair Repulsion model is used 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 of bond 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

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Shapes that minimize repulsions:

linear triangular planar

tetrahedral triangular bipyramidal

octahedral

Predicting Molecular Shapes: VSEPR

Bonds and lone pairs determine shape.

Use the notation AX_nE_m n atoms bonded to

central atom A

m lone pairs on central atom A

If a molecule contains:

• bonding pairs only – these angles are correct:

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

• lone pair/lone pair repulsions are largest.

• lone pair/bond pair are intermediate in strength.

• bond/bond interactions are the smallest.

Predicting Molecular Shapes: VSEPR

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.

Predicting Molecular Shapes: VSEPR

2 e^-groups bond lone pairs pairs

2 0 AX₂E₀

linear

....

1 1 AX1E1

linear

Linear e^-pair geometry

molecular geometry

Predicting Molecular Shapes: VSEPR

AX_nE_m: 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.

O C O

Cl Be Cl

Each H-C-C unit is linear.

H C C H

180.0°

Predicting Molecular Shapes: VSEPR

AX₂E₀examples:

3 e^-groups bond lone pairs pairs

.... ....

3 0 AX3E0

triangular planar

2 1 AX₂E₁

angular (bent)

1 2 AX1E2

linear

Triangular planar e^-pair geometry

molecular geometry

Predicting Molecular Shapes: VSEPR

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AX₃E₀examples:

Triangular planar.

Each C is AX₃E₀= triangular planar.

Cl B Cl

Cl

C

H C H

H H

120°

Predicting Molecular Shapes: VSEPR

4 e^-groups

bond lone

pairs pairs

4 0 AX₄E₀

tetrahedral

.. ..

AX₁E₃?

All molecules with only 1 bond are linear!

3 1 AX3E1

triangular pyramidal

2 2 AX2E2

angular (bent)

.. ..

Predicting Molecular Shapes: VSEPR

Tetrahedral e^-pair geometry

molecular geometry

AX₄E₀

All angles = tetrahedral angle.

AX₃E₁

Lone-pair/bond > bond/bond repulsion: H-N-H angle is reduced.

AX2E2

Two lone pairs: H-O-H angle is even smaller.

H C H

H H

H N H

H

O H

H

Predicting Molecular Shapes: VSEPR

VSEPR applies to each atom in a molecule.

• Alkanes: each C is tetrahedral.

Predicting Molecular Shapes: VSEPR

Tetrahedral O

Lactic acid:

Tetrahedral C

Triangular planar C

Tetrahedral C H

C H C

H C O

O O H

H H

..

.. .. ..

.. ..

Tetrahedral O

Predicting Molecular Shapes: VSEPR

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

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The atoms are non-equivalent.

Green atoms are axialaxial; blue atoms are equatorialequatorial.

Expanded Octets

AX_nE_m: m + n = 5

Triangular bipyramidal e^-pair geometry.

Triangular

bipyramidal Seesaw T-shaped Linear

F P F

F

F F

F S F

F F

F Cl F

F

F Xe F

Expanded Octets

AX_nE_m: m + n = 6 Octahedral e^-pair geometry:

Octahedral Square pyramid Square planar All atoms are

equivalent in

AX₆E₀ F S F

F

F F

F

F Br F

F

F F

Cl I Cl

Cl Cl

Expanded Octets

Lewis dot + VSEPR predict molecular shapes, butbut…

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

Valence bond theory

Valence 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 H₂and HF, but why does…

• Be form compounds?

• Be (1s²2s²).

• No unpaired e^-to share.

• Experiments show: linear BeH₂, BeCl₂, …

• C form 4 bonds at tetrahedral angles?

• C (1s²2s²2p²).

• 2p_x¹2p_y¹ Two bonds?

• p orbitals are at 90° to each other

• Experiments show: tetrahedral CH4, CCl4, …

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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 porbital → two sp hybrids.

Orbitals Consistent with Molecular Shapes

sp Hybrid Orbitals

Be compounds (BeH₂, BeF₂…):

Each sp hybrid (180° apart) holds one e^-. Two equivalent covalent bonds form.

sp

²

Hybrid Orbitals

B forms three sp²hybrid orbitals:

• One s orbital mixes with two p orbitals.

• One p orbital remains unmixed.

sp

²

Hybrid Orbitals

B compounds (BH₃, BF₃…):

Each sp²hybrid (120° apart) holds one e^-. Three equivalent covalent bonds form.

sp

³

Hybrid Orbitals

C forms four sp³hybrid orbitals:

• One s orbital mixes with three p orbitals.

• All p orbitals are mixed.

In C, each sp³hybrid (109.5° apart) holds one e^-. Four equivalent covalent bonds form.

sp

³

Hybrid Orbitals

N and O compounds (NH₃, H₂O…) have more e^-:

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sp

³

Hybrid Orbitals

“Octet rule” molecules have tetrahedral e^-pair shape.

• sp³hybridized (CH₄, NH₃, H₂O, H₂S, PH₃, …)

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

There are:

• 4σ bonds in CH₄

• 3σ bonds in NH₃

• 2σ bonds in H₂O

H

C

H H 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 C

H C H

H H

The double bond in ethene is composed of:

• aσσ bondbond – head-to-head overlap of sp²on each C atom.

• aππ 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 bond

C C

H H

H overlap

Sideways overlap forms oneoneπ 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:

• oneσσ bondbond

• twoππ bondsbonds

• sp hybridization leaves two unmixed p orbitals on each C.

C C

H H

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σ bond: C (sp) + C (sp) overlap:

C C H

H

Two

Twoπ bonds

• above and below overlaps are 1 bond.

• front and back overlaps are a second bond.

Two

Two p orbitals on eacheach C contain a single e^-.

C C H

H overlap H C C 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 isomerism to occur.

• Most bonds are polar (e.g. C-O)

• O isδ-, C is δ+ (EN_O= 3.5, EN_C= 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 toδ-, the + showsδ+

O = C = O

δ-

δ- 2δ+

Molecular Polarity

• Water is polar (bond dipoles do not cancel) Dipole,μ = 1.85 D

H H O

+ Net dipole

Molecular Polarity

Dipole moment

Dipole moment (μ) is a measure of molecule 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 0

nonpolar (μ=0) highly polar weakly polar

A molecule is nonpolarnonpolar if it is:

• AX_nEE₀₀and all XX are identical.

CO₂ AX₂E₀ linear CH₄ AX₄E₀ tetrahedral CCl₄ AX₄E₀ tetrahedral

PF₅ AX₅E₀ triangular bipyramidal

• “divisible” into nonpolar AX_nE₀shapes

PCl₃F₂ triangular planar (PCl₃) + linear (PF₂) XeF₄ linear (XeF₂) + linear (XeF₂)

Molecular Polarity

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AX_nE_mmolecules are polarpolar if they don’t divide into nonpolar shapes, and::

Molecular Polarity

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

• m≠ 0:

H₂O AX₂E₂ bent polarpolar NH₃ AX₃E₁ pyramidal polarpolar

• The X in AX_nE₀differ:

CH₂Cl₂ AX₄E₀ tetrahedral polarpolar PF₄Cl AX₅E₀ triangular bipyramidal polarpolar

Molecular Polarity

F F

F

C F

CF₄is non polar No net dipole

F F

H

C F

CHF₃is polar Net dipole +

Non polar Non polar AX₅E₀; identical X

PCl₅

PCl4F

Non polar Non polar AX5E0and “X” differ.

BUT divisible into nonpolar shapes:

linear + triangular linear + triangular

planar planar

PF₃Cl₂ Polar

Polar AX₅E₀

“X” differ +

Polar Polar AX5E0and “X” differ.

Doesn’t divide into nonpolar shapes

Molecular Polarity

PCl3F2

Molecules attract each other.

Intermolecular Intermolecular 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 Forces

Also called dispersiondispersion forces.

• Random e^-motion produces a temporary dipole in one 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

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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 is hard 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 largeδ+ and δ- values.

H-bonds are usually drawn as dotted lines.

Hydrogen Bonding

H on one molecule interacts with O on another molecule.

Hydrogen Bonding

Water is a liquid at room T (not a gas).