KLEIN Chap 1 Handout

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1

In the mid 1800s, it was first suggested that substances

are defined by a specific arrangement of atoms.

Why is a compounds formula NOT adequate to define it?

What term do we use to describe different substances

with the same formula?

Chapter 1 Structural Theory

Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-1

Atoms that are most commonly bonded to carbon

include N, O, H, and halides (F, Cl, Br, I).

With some exceptions, each element generally forms a

specific number of bonds with other atoms:

Practice with SKILLBUILDER 1.1.

1.2 Structural Theory


How do potential energy and stability relate?

What forces keep the bond at the optimal

length?

1.3 Covalent Bonding


For simple Lewis structures:

1. Draw the individual atoms using dots to represent the

valence electrons.

2. Put the atoms together so they share PAIRS of electrons to

make complete octets. WHAT is an octet?

Take NH3, for example:


1.3 Covalent Bonding Simple Lewis

Structures


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Consider the formal charge example below. Calculate the formal

charge on each atom.

Carbon SHOULD have 4 valence electrons because it is in Group IVA on the

periodic table.

Carbon ACTUALLY has 8 valence electrons. It needs 8 for its octet, but only

4 count towards its charge. WHY?

The 4 it ACTUALLY has balance out the 4 it SHOULD have, so it does not

have a formal charge. Its neutral.

1.4 Formal Charge


or

Covalent bonds are either polar or nonpolar:

Nonpolar covalent bonds: bonded atoms share electrons evenly

Polar covalent bonds: one of the atoms attracts electrons more than the

other

Electronegativity:

How strongly an atom attracts shared electrons

1.5 Polar Covalent Bonds


Electrons tend to shift away from lower

electronegativity atoms to higher electronegativity

atoms.

The greater the difference in electronegativity, the more

polar the bond.

1.5 Polar Covalent Bonds


Electrons behave as BOTH particles and waves. How can they be

both?

Maybe the theory is not yet complete.

The theory does match experimental data, and it has predictive

capability.

Like a wave on a lake, an electrons wavefunction can be (+), (-), or ZERO.

1.6 Atomic Orbitals


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Because they are generated

mathematically from wavefunctions,

orbital regions can also be (-), (+), or

ZERO.

The sign of the wave function has nothing

to do with electrical charge.

In this p-orbital, there is a nodal plane.

The sign of the wavefunction will be

important when we look at orbital

overlapping in bonds.

1.6 Atomic Orbitals


Once the 2s is full, electrons fill into the three

DEGENERATE 2p orbitals.

Where are the nodes in each of the 2p orbitals?

1.6 Atomic Orbitals


The bond for an H2 molecule results from constructive

interference.

Where do the bonded electrons spend most of their

time?

1.7 Valence Bond Theory


Atomic orbital wave functions

overlap to form molecular

orbitals (MOs) that extend

over the entire molecule.

MOs are a more complete

analysis of bonds because they

include both constructive and

destructive interference.

The number of MOs created

must be equal to the number

of atomic orbitals that were

used.

1.8 Molecular Orbital Theory


H2 MOs

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Imagine a He2 molecule.

How would its MOs

compare to those for H2?

How would the energy of

the He2 compare to 2 He?

Why does helium exist in its

atomic form rather than in

molecular form?

In general, if a molecule has

all of it MOs occupied, will it

be stable or unstable?

1.8 Molecular Orbital Theory


H2 MOs

To make CH4, the 1s atomic orbitals of the H atoms will

overlap with the four sp3 hybrid atomic orbitals of C.

1.9 Hybridized Atomic Orbitals


Consider ethene (ethylene).

Each carbon in ethene must bond to THREE other

atoms, so only THREE hybridized atomic orbitals are

needed.



An sp2 hybridized carbon will have three equal-energy

sp2 orbitals and one unhybridized p orbital.



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The sp2 atomic orbitals overlap to form sigma () bonds.

Sigma bonds provide maximum HEAD-ON

overlap.



The unhybridized p orbitals in ethene form pi () bonds,

created by SIDE-BY-SIDE orbital overlap.

Practice with CONCEPTUAL CHECKPOINT 1.20.



The sp atomic orbitals overlap HEAD-ON to form sigma

() bonds while the unhybridized p orbitals overlap

SIDE-BY-SIDE to form pi () bonds.




Explain the different strengths and lengths below.

Practice with CONCEPTUAL CHECKPOINT 1.24.



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Valence shell electron pair repulsion (VSEPR theory):

Valence electrons (bonded and lone pairs) repel each other.

To determine molecular geometry:

1. Determine the steric number.

1.10 Molecular Geometry


To determine molecular geometry:

2. Predict the hybridization of the central atom:

If the steric number is 4, then it is sp3.

If the steric number is 3, then it is sp2.

If the steric number is 2, then it is sp.



For any sp3 hybridized atom, the four valence electron pairs

will form a TETRAHEDRAL ELECTRON GROUP geometry:

Methane has four

equal bonds, so the

bond angles are

equal.

HOW does the

lone pair of ammonia

affect its geometry?

The bond angles in

oxygen are even

smaller. WHY?


sp3 Geometry


The MOLECULAR geometry is different from the

ELECTRON GROUP geometry. HOW?

ExampleSteric

numberHybridization

Arrangement

of electron

pairs

Arrangement of

atoms

(geometry)

CH4 4 sp3 Tetrahedral Tetrahedral

NH3 4 sp3 Tetrahedral Trigonal pyramidal

H2O 4 sp3 Tetrahedral Bent


sp3 Geometry


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Calculate the steric number for BF3.

Electron pairs that are located in sp2 hybridized orbitals

will form a trigonal planar ELECTRON GROUP geometry.

What will be the molecular geometry?


sp2 Geometry

Copyright 2012 John Wiley & Sons, Inc. Klein, Organic Chemistry 1e1-25 Practice with SKILLBUILDER 1.8.


Summary


Electronegativity differences cause induction.

Induction (shifting of electrons WITHIN their orbitals)

results in a dipole moment.

Dipole moment = (the amount of partial charge) x (the

distance the + and - are separated)

Dipole moments are reported in units of debye (D)

1 debye = 10-18 esu cm

An electrostatic unit of charge (esu) is a unit of charge. One electron

has a charge of 4.80 x 10-10 esu.

Centimeters (cm) are included in the unit because the distance

between the centers of + and charges affects the dipole.

1.11 Molecular Polarity


Consider the dipole for CH3Cl

Dipole moment () = charge (e) x distance

(d)

Plug in the charge and distance

= (1.056 x 10-10 esu) x (1.772 x 10-8 cm)

Note that the amount of charge separation is

less than what it would be if it were a full

charge separation (4.80 x 10-10 esu).

= 1.87 x 10-18 esu cm

Convert to debye

= 1.87 D



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Check out the polarity of come other common bonds:



For molecules with multiple polar bonds, the dipole

moment is the vector sum of all of the individual bond

dipoles.



Electrostatic potential maps are often used to give a

visual depiction of polarity.





Klein, Organic Chemistry 1eCopyright 2012 John Wiley & Sons, Inc. 1-32

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Dipoledipole forces result when polar molecules line

up their OPPOSITE charges.

Note that acetones permanent dipole results from the

difference in electronegativity between C and O.

The dipoledipole attractions BETWEEN acetone

molecules affects acetones boiling point (BP) and

melting point (MP). HOW?

1.12 Intermolecular Forces

Dipoledipole


Why do isobutylene and acetone have such different

MPs and BPs?


Dipoledipole


Hydrogen bonds (H-bonds) are an especially strong type

of dipoledipole attraction.

Hydrogen bonds are strong because the partial + and

charges are relatively large.

Why are the partial charges in the H-bonding examples

below relatively large?


Hydrogen Bonding


Explain why the following isomers have different boiling

points.


Hydrogen Bonding


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The result is a fleeting attraction between the two

molecules.

Such fleeting attractions are generally weak.

However, like any weak attraction, if there are enough

of them, they can add up to a lot.


London Dispersion Forces


Explain why molecules with more mass generally have

higher boiling points.





Explain why more highly branched molecules generally

have lower boiling points.




We use the principle that like dissolves like.

Polar compounds GENERALLY mix well with other polar

compounds:

If the compounds mixing are all capable of H-bonding and/or

strong dipoledipole interactions, then there is no reason why

they shouldnt mix.

Nonpolar compounds GENERALLY mix well with other

nonpolar compounds:

If none of the compounds are capable of forming strong

attractions, then no strong attractions would have to be

broken to allow them to mix.

1.13 Solubility


KLEIN Chap 1 Handout

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