Atomic Structure and Bonding

–Covalent bonds are formed by the overlap of two atomic orbitals and the electron pair is shared by both atoms. ... ΔX > 1.9 ionic bond ΔX < 0.5 coval...

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Atomic Structure and Bonding

Chapter 1 Organic Chemistry, 8th Edition John McMurry

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Common Elements Groups

First row Second row

In most organic molecules carbon is combined with relatively few elements

Columns 2

Lewis’ Model • In Lewis structures electrons are represented as dots. • Three general rules. – Include only valence electrons. – If possible, every 2nd raw element should have 8 electrons. – Every H atom has 2 electrons.

2 electrons (He)

8 electrons (Ne)

Lewis structure

Kekulé structure

3

Molecular Shape • The molecular structure is defined by: – bond lengths – bond angles Bond lengths decrease along a period.

bond length Bond lengths increase along a group

Bond length 4

Bond Lengths

Bond

Length (Å)

Bond

Length (Å)

Bond

Length (Å)

5

Geometry – VSEPR Theory • The number of Valence Shell Electron Pairs (groups) around an atom defines the geometry of that atom. • A group is an atom or a non bonding pair of electrons. • Groups will tend to be as far apart as possible.

Number of groups

Geometry

Angle

2

linear

180°

3

trigonal planar

120°

4

tetrahedral

109.5° 6

4 groups: CH4 Tetrahedral

Square planar

Preferred

Not observed

Tridimensional representations of methane

7

3 groups: BF3 and C2H4 2 trigonal molecules

ball-and-stick model

3 atoms around B

3 atoms around each C

All 3 atoms are in the plane

All 6 atoms are in the plane space-filling model

8

2 groups: BeH2 and C2H2

2 linear molecules



2 atoms around Be

2 atoms around each C

9

Energies of Multiple Bonds Bond

Bond Dissociation Energy (kJ/M) 360 700 950 400 750 360 700

950

The Lewis model is not adequate!

Models for the Chemical Bond • Valence Orbital theory. – Covalent bonds are formed by the overlap of two atomic orbitals and the electron pair is shared by both atoms. – A valence bond is localized between two atoms.

• Molecular Orbital theory. – n atomic orbitals are combined to give a new set of n molecular orbitals (bonding and antibonding). – Molecular orbitals are delocalized on the whole molecule.

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Valence Orbitals • Bonds are formed by the in-phase overlap of two atomic orbitals each contributing one electron. • The electron pair is localized between two atoms and is shared by both atoms. • Hydrogen uses the 1s orbital to form s bonds. • 2nd row atoms use hybrid orbitals (sp3, sp2, sp) to form s bonds. • 2nd row atoms use p orbitals to form p bonds that have a nodal plane. • Atomic orbitals overlap better in s bonds (co-linear) than in p bonds (parallel).

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Strong Bonding Interactions: Axial Overlap X 1s(H)-1s(H) H2

X 2px(F)-1s(H) F-H

X 2px(F)-2px(F) F-F

Weak Bonding Interactions: Lateral Overlap Z

Z

2pz(C)2pz(C) 2pz(C)2pz(C) liaison s of ducyclopropane cyclopropane s bond liaison l'éthylene p bondpofde ethylene

Inexistent Bonding Interactions (0 Overlap) Z

Z X

2px(F)-1s(H) F-H

y

2pz(C)2pz(C) liaisonintetra-tBu-ethylene Bond tetra-tBu-ethylene

X

Molecular Orbitals • Valence electrons occupy molecular orbitals delocalized on the whole molecule. • The combination of n atomic orbitals gives n new molecular orbitals. • Bonding orbitals have lower energies and antibonding orbitals have higher energies than the starting atomic orbitals

The H2 molecule:

Robert Mulliken (1896-1986)

DE

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Atomic Orbitals of Carbon

s

p

2p

3 x 2p orbitals

sp3 Hybrids • The mixing of a spherical 2s orbital and three 2p orbitals generates four sp3 orbitals, each with a small and a large lobe.

hybridation

2s

3 x 2p

4 x sp3 orbitals

p

sp3

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sp3 Hybrids

2p sp3

2s

E

Methane The overlap of a half-full 1s orbital of hydrogen with a half-full sp3 orbital of carbon bond gives a s orbital.

+ s H

s C–H

+

H

+ C–

C–

sp3

s s

s s 19

Ethane

ethane

tetrahedral sp3 C Two sp3 hybrids overlap giving the C–C s bond

sp3 hybrids on C overlap with 1s orbitals on H giving the CH s bonds.

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Sp2 Hybrids

2p

p sp2

2s

E

Ethylene C2H4

3 groups around C C atoms are sp2

Sp2 hybrids

C-C double bond

Sp Hybrids

2p

p

sp

2s

E

Acetylene C2H2

C-C triple bond

Structures of C2H6, C2H4, C2H2

109.6°

dC-C (pm): 154 dC-H (pm): 110 EC-C (kJ/M): 376

121.7°

133 107.6 611

180°

120 106 835

Energy Levels and Orbital Size E

2p

2sp3 2sp2 2sp

2s

C

Carbone

C*

Carbone activé

C sp3

3

Carbone hybridé sp

C sp2

2

Carbone hybridé sp

C sp

Carbone hybridé sp

Ibridization of O, S, N

[He]2s22p4

[He]2s22p3

[Ne]3s23p4

Polar Bonds Intermolecular Interactions Delocalised Bonds Chapter 2 th Organic Chemistry, 8 Edition John E. McMurry

Polar Covalent Bonds • The higher the electronegativy difference, the higher the polar character of a covalent bond. δ+ δ-

δ+ δ-

δ+

δ- δ+ δ-

• In polar bonds, bonding electrons are attracted towards the more electronegative atom.

• Generally:  ΔX > 1.9 ionic bond  ΔX < 0.5 covalent bond  ΔX = 0.5 – 1.9 polar covalent bond

Pauling’s Electronegativities

Intermolecular Interactions 

Intermolecular interactions are also called non-covalent and non-

bonded interactions.



Intermolecular interactions depend on the type and number of functional groups. In neutral molecules there are three main types of intermolecular interactions. • Vand der Waals interactions (London dispersion forces) – VDW • Dipole-dipole interactions – DD • Hydrogen bonds– HB

strenth



Dipole Moments •

Polar molecules have one or more polar bonds. Es. H2O



Apolar molecules either do not have polar bonds or have polar bonds whose dipoles cancell each other. E.g. CO2

CCl4 d = 0 D

Dipoles cancel

CH2Cl2 d = 1.62 D

Dipoles add

The Hydrogen Bond The hydrogen bond is an electrostatic interaction between a O-H or N-H group and a lone pair on O or N.

Hydrogen bond

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Dipole-Dipole Interactions Dipole-dipole interactions are attraction forces between the permanent dipoles of two molecules.

Dipole-dipole interactions

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Van der Waals (London) Forces 

VdW forces are weak interactions originating from temporary variations of the molecule’s electron density distribution.



They are the only attractive forces in apolar molecules. Van der Waals interactions between two CH4 molecules

Dipoles generated by a temporay asymmetry in the electron density

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Van der Waals (London) Forces 

Van der Waals interactions are present in all molecules.



The larger the surface area, the larger the attractive force between two molecules, and the stronger the intermolecular forces. Long, cylindrical molecules: stronger interactions

Compact, spherical molecules: weaker interactions

n-pentane

neopentane

Van der Waals (London) Forces  

VdW forces depend on polarizability. Larger atoms, like iodine, which have more loosely held valence electrons, are more polarizable than smaller atoms like fluorine, which have more tightly held electrons Small atoms: lower polarizability Weak interaction

Large atoms: higher polarizability

Stronger interaction

Summary Interaction Van der Waals

VDW Dipole-dipole DD Hydrogen bond HB ionic

Relative strength

Present in

Very weak

All molecules

weak

Permanent dipoles

strong

Very strong

Molecules with OH, NH, FH funct. groups Ionic compounds

Polar molecules interact strongly than apolar ones.

Examples

Boiling Point 

The boiling point is the temperature at which the vapor pressure of a liquid is equal to the external pressure.



Energy is required to break intermolecular interactions.



The higher the intermolecular interactions, the higher the b.p..



Compounds with similar M.W.:

Van der Waals

Dipole-dipole

Hydrogen bond

Boiling point

pentane (m.w. 72)

butanal (m.w. 72)

1-butanol (m.w. 74)

Boiling Point

b.p. = 42 °C I is more polarizable

b.p. = 102 °C Larger surface area

b.p. = –78 °C Smaller F has a low polarizability

b.p. = 56 °C Smaller surface area

Melting Point 

M.p. and b.p. follow the same trend.

pentane

butanal

Melting point

1-butanol

Delocalized Electrons and Bonds. Resonance • The structure and properties of certain molecules can not be explained by the simple valence orbital model with localized electrons.

• In this case, a single Lewis structure is replaced by a set of Lewis structures: the molecule is said to resonate between these structures and this phenomenon is called resonance. CH3CH2OH: pKa 16

CH3COOH: pKa 4.75

Resonance structures Localized charge less stabile

Delocalized charge more stabile

Resonance hybrid

Resonance • Resonance structures have the same disposition of atoms but a different arrangement of electrons (π electrons and lone pairs).

• Bond lengths and angles do not change in resonance structures. • Resonance is a simple theory to adapt conventional Lewis structures to the representation of molecules with delocalized electrons and bonds.

Resonance Benzene (C6H6)

Very stabile 6 identical C-C bonds

Resonance structures

E N E R G Y

Resonance hybrid

6 localized p electrons Er = 150.7 kJ/mol (36 kcal/mol) 6 delocalized p electrons

Resonance

Resonance

1. Resonance structures are not real. No single resonance structure can adequately represent the real structure of a species with delocalized electrons. 2. Resonance structures are not isomers. They only differ in the distribution of electrons not in the disposition of the nuclei. 3. Resonance structures are not in equilibrium.