1. Lecture 1 - Structure and Bonding I PDF

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CHEMISTRY 138 – Organic Chemistry I - a course for students of the life sciences

Questions? Ask as we go!

Objective: to learn some of the fundamental principles which allow us to understand the reactions of organic molecules. Structure and bonding

Reactivity of some important functional groups

Aromaticity

An introduction to spectroscopy (using infra-red light to detect functional groups)

Structure and Bonding Chapter 1, McMurry 8th Edition

Structure: the shapes of organic chemistry C 2H 2

C 2H 4

CH4

linear

planar

tetrahedral

These shapes recur in all organic molecules:

E.g.

Structure: drawing organic molecules

H

H C

Dashed wedges indicate a bond sticking back behind the plane.

H H

C

H

C

H

H

H Wedges indicate a bond sticking out in front of the plane.

Lines indicate a bond in the plane.

Other common representations:

H N H

H

H

H

H

H

C

C

C

C

C

H

H

H

H

H

H N H

Each vertex is a carbon. Each "unstated bond" is a hydrogen.

O O

HO

Zingerone

From structure to bonding Structure of the atom  Schrödinger model: atomic orbitals

Bonding in organic molecules  covalent bonds  ionic bonds

valence bond theory Lewis structures: convenient bonding description

VSEPR: deriving the shapes of organic chemistry

hybridization: explaining the shapes of organic chemistry

Electronic structure of the atom: Schrödinger model (1926) de Broglie: matter with wave properties wave equation:

HYEY



h mv

- has many solutions

Y = "wavefunction" or "orbital" = a mathematical function having wave characteristics. - describes volume of space where electron is MOST LIKELY to be found

Solving H Y  E Y for H-atom, 1s orbital:

2s orbital:

Solving H Y  E Y for H-atom, 2p orbitals:

3d orbitals:third row elements (e.g., S, P, Cl)

node

From

H Y  E Y,

many solutions

Y = "wavefunction" or "orbital" E = energy of each orbital

For the H-atom, shell

E shell

shell

Theory predicts that, for the H-atom, the s, p, and d orbitals are degenerate, i.e. will have the same energy.

Theory predicts that, for any atom other than H, the s, p, and d orbitals will have different energies:

shell

Rules for filling orbitals with electrons:

1. Lowest energy orbitals fill first 2. 2 electrons per orbital with opposing spin 3. Electrons space themselves out before pairing (if two or more orbitals are available with equal energy, put one electron in each until all the orbitals are half full.)

Theory predicts that, for any other atom other than H, the s, p,

and d orbitals will have different energies.

shell

H Li Be

He B

C

N

O

Na

BONDING Bonds form because the molecule formed has a lower energy than the atoms apart. Generally, electrons are shared or transferred so that each atom has a filled valence shell.

Ionic bonding: "Perhaps one of you gentlemen could tell me just what it is outside the window that you find so attractive...?"

F

Ne

Ionic bonding: electrostatic forces between oppositely charged ions.

ionic bond

filled valence shells

Li+ +

F



Li+ F



Ionic compounds form because opposite charges attract. Coulombic attraction provide enough energy to cause salt formation.

Covalent bonding: electrons shared between atoms. unpaired electron, unfilled valence shell

filled valence shells for both

Energy

H H

H H

0 HH

Distance

H

H

The smaller the separation, the stronger the bond... …until nuclear-nuclear repulsions become too large.

1. Valence Bond (VB) Theory 2. Molecular Orbital (MO) Theory Valence Bond (VB) Theory Covalent bonds: two electrons shared between two nuclei = "localized" covalent bonds Lewis structures: a representation of covalently bonded molecules = covalently bonding pair •• = non-bonding pair

H C H H H

Each atom is surrounded by a filled valence shell of electrons. "non-bonding" electron pairs

Some practice drawing Lewis structures:

Valence Shell Electron-Pair Repulsion Model = VSEPR -an understanding of the shapes of molecules in terms of electrostatic repuls'n between electron groups surrounding the atom electron group = bonding pair, multiple bonding group or lone pair of electrons

H

C

O

N

H

O

H F F

C

N

F

From the valence atomic orbitals of carbon…

2s

2px

C

2p __ __ __ 2s __

2py

how can we explain this?

The concept of hybridization is needed to explain the geometry of molecules.

2pz

Valence Bond (VB) Theory:

H H

C

H H methane has four identical bonds

unhybridized

tetrahedral

sp3 hybridization: s+3p

sp3 hybridization: s+3p

tetrahedral

● ●

H

● ●

H

●

●

C

H

sp3 hybridized

H  (sigma) bond: lies along an imaginary axis between the two nuclei of that bond

Valence Bond (VB) Theory:

H

H C

C

H

H

unhybridized trigonal planar

sp2 hybridization: s+2p

C

2p

H

•

•

C

C

H

+

●

H H

●

C

●

C

● ● ● ●

H H

●● ● ●

●

sp2 hybridization: s+2p

C

C

C

2p

H H trigonal planar

H

H C

H

C H

π bond does not lie along axis between nuclei

sp2 hybridized

Valence Bond (VB) Theory:

unhybridized

linear

sp hybridization: s+p

H

C

C

O

H

Ethane 109.5o

Ethylene

H

bond angle

120o

C

C-C bond length

C-C bond energy

154 pm

376 kJ/mol

O

Acetylene 180o

H

H

C

H

C

C

C

H

133 pm

611 kJ/mol

120 pm

835 kJ/mol

A double bond is less than twice as strong as a single bond.

H

Polar Bonds In covalent bonds, electrons are not always shared equally.

H

H

Non-polar covalent bond (equal sharing)

H

F

Polar covalent bond

δ+

H

δ-

H

F

F

partial partial positive negative charge charge

Polar bonds

Electronegativity of an atom: the ability of an atom in a molecule to attract electrons toward itself.

bond dipole









 H 3C

 MgBr

In organic molecules, atoms with high electronegativities are said to be electron withdrawing.

ΔEN = 0

ΔEN < ~2.0

ΔEN > ~2.0

H2 HF LiF

EN (Cl) = 3.0

EN (C) = 2.5

EN (C) = 2.5

EN (H) = 2.1

ΔEN = 0.5

ΔEN = 0.4

Dipole Moment – measurable quantity – describes the separation of charge in a molecule

- charge separation: O is  – both H are  +

Dipole moment of the molecule is the vector sum of all of the bond dipoles

Which of the following molecules have a dipole moment? If it has one, in which direction does it lie?

H C Cl

Cl Cl

Formal Charge – which atoms in the molecule bear formal charge? Some examples: Periodic Table

H3

O+ :

H O H H

1

4

5

6

7

H Li

C

N P

O S

F Cl Br I

bromonium cation:

+

nitromethane:

+ _ Sum of all formal charges must equal the charge on the molecule

Resonance

154 pm

133 pm

How does this arise?

+

•

How can we describe this bonding situation?

•

•

•

+

Why does this arise? The bonding described by the combination of resonance forms is more stable than the bonding described by any individual resonance form.

The two resonance forms of benzene

•

• •

The “real structure” is a mixture of these two resonance forms

•

•

•

In drawing resonance forms, the arrangement of the electrons changes but the nuclei do not move. Resonance forms differ only in the placement of the -electrons and the non-bonding electrons (lone pairs).

O(-)

O

Individual resonance forms do not have to be equivalent.

C CH3

O

(-)

CH3

C

equivalent resonance forms

Resonance forms must be valid Lewis structures that obey the rules of valence.

-

-

Using arrows to show 'flow' of electrons between resonance forms +

•

•

•

•

-

+

The doublesided arrow indicates the direction of movement of a pair of e-

O

Criteria for determining which resonance forms are the most important (in decreasing importance): 1. Number of -bonds. Since it costs energy to break bonds, the resonance structures with the most -bonds make the greatest contribution to the stability of a species. OCH 3

+ CH2

2. Formal charge. (a) Resonance structures with greater charge separation are less important than those with lesser.

O

O

C H 3C

C N H

R

H3 C

R

N H

(b) In resonance forms with equal formal charges, the more important forms have the negative charge on the more electronegative atom. O

O C H 3C

-

CH 3

C

CH3 H 3C

C H

C H

3. Inductive effects. Electron donating groups help to stabilize (+) charges. Electron withdrawing groups help to stabilize (-) charges and destabilize (+) charges.

CH3

CH3

CH3 H

H

H

+ +

NO2

+

NO2

NO2

Drawing correct resonance forms takes practice!

A very important idea: From molecular structure comes everything else. Functional groups: a group of atoms within a molecule that has characteristic chemical behaviour.

O

vs.

- a large number of functional groups in organic chemistry - be able to recognize and name important functional groups (see next page)

Structures of Some Important Functional Groups alkene

alkyne OH

alcohol

ether

alkyl halide (Cl, Br, I)

amine

CH3

O

CH(CH3)2

CH3CH2

N

CH2CH3

H

Br

O

O

ketone aldehyde

O

ester

O

O O

carboxylic acid

OH

amide

N H...