# Review

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29 Οκτ 2013 (πριν από 4 χρόνια και 6 μήνες)

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Organic Chemistry

Review Information for Unit 1

Atomic Structure

MO Theory

Chemical Bonds

Atomic Structure

Atoms are the smallest representative particle
of an element.

Three subatomic particles:

protons

neutrons

electrons

All atoms of a particular element have the
same number of protons and electrons.

The number of neutrons can vary.

nucleus

Atomic Structure

Isotope

atoms with the same number of protons but
different numbers of neutrons

Carbon has three isotopes:

12
C:

6 protons, 6 neutrons

most common isotope

13
C:

6 protons, 7 neutrons

used in structural determinations

14
C:

6 protons, 8 neutrons

used to determine the age of organic
materials (C
-
14 dating)

Atomic Structure

The location and energy of the electrons in an
atom are best described using the Quantum
Mechanical Model.

Electrons have both particle
-
like and wave
-
like properties.

Solving the Schroedinger equation leads to a
series of mathematical functions called wave
functions (
y
)

describe an allowed energy state for an
electron

(
orbital
)

Atomic Structure

Heisenberg Uncertainty Principle
:

the exact energy and exact location of an
electron in an atom cannot be known
simultaneously

Since solutions to the Schroedinger equation
give the exact energy of the electron,
the
exact location of the electron is uncertain.

Electrons don’t move in well
-
defined circular
orbits around the nucleus
.

Atomic Structure

Although we cannot determine the exact
location of an electron, an orbital describes a
specific distribution of
electron density

in
space

the probability of finding an electron in a
particular region of space

The QM model uses 4 quantum numbers to
describe an electron in an orbital
:

n, l, m
l

are used to describe the orbital
itself

m
s

is used to describe the spin of the
electron.

Atomic Structure

Quantum Numbers:

principal quantum number (n)

energy of the electron

relative distance from the nucleus

azimuthal quantum number (l)

shape of the orbital

magnetic quantum number (m
l
)

orientation in space of the orbital

electron spin quantum number (m
s
)

direction of electron spin

Atomic Structure

There are four common “types” of orbitals

s orbital

spherical

One per subshell

p orbitals

3
-
D figure 8

3 per subshell when n
>
2

same energy

different orientation in space

1s

z

y

x

p
x

Atomic Structure

x

x

y

z

z

y

The p
x
, p
y
, and p
z

orbitals are superimposed
at 90
o
angles.

The p
x

and p
y

orbitals are in the plane of the
slide while the p
z

orbital comes out toward
you at 90
o

from the plane of the slide.

Atomic Structure

d orbitals

5 per subshell when n
>
3

same energy

different orientations in space

complex shapes

f orbitals

7 per subshell when n
>
4

same energy

different orientations in space

complex shapes

Note:

Most organic
compounds do not utilize
d and f orbitals.

Atomic Structure

An
orbital diagram

or the

electron
configuration

can be used to describe the
arrangement of the electrons in the orbitals of
an atom.

According to the
aufbau principle,

the
electrons in an atom in the ground state will be
found in the lowest energy orbital that is
available.

Atomic Structure

Use the diagonal diagram to determine the
relative energies of the orbitals:

1s

2s

2p

3s

3p

3d

4s

4p

4d

4f

5s

5p

5d

5f

Atomic Structure

The
Pauli Exclusion Principle

and
Hund’s rule

also govern the placement of electrons.

Pauli Exclusion Principle:

No two electrons in an atom can have the
same set of four quantum numbers n,
l
, m
l
,
and m
s
.

Maximum of 2 electrons with opposite
spins per orbital

Atomic Structure

Hund’s Rule
:

If more than one orbital with the same
energy is available, electrons will fill empty
orbitals first.

Keep electrons unpaired as long as an
empty orbital with the same energy is
available.

Atomic Structure

Example:

Draw an orbital diagram and write the
electron configuration of N.

# electrons =
7

1s

2s

2p

Orbital diagram:

Atomic Structure

# electrons =
7

Electron configuration:
1s
2
2s
2
2p
3

This electron configuration does not clearly
indicate that all three 2p electrons are unpaired.

A better representation is to clearly show where
each p electron is found:

1s
2
2s
2
2p
x
1
2p
y
1
2p
z
1

Atomic Structure

On your exam, you should be able to draw an
orbital diagram or write electron configurations
that clearly indicate the location of each
electron.

i.e. show whether an electron is in a p
x
, p
y
,
or p
z

orbital.

MO Theory

Quantum mechanics describes the electrons in
an
atom

using wave functions called
atomic
orbitals.

Allowed energy states for electrons in an
atom in the QM model

According to Molecular Orbital (MO) Theory, a
chemical bond

is formed when 2 atomic orbitals
on different atoms

overlap and combine

Two new
molecular orbitals

are formed:

bonding molecular orbital

antibonding molecular orbital

MO Theory

Molecular orbitals formed when two hydrogen
atoms combine:

Antibonding
molecular orbital

Bonding molecular
orbital

Bonding molecular orbital

Constructive interference between to atomic
orbitals leads to a build up of e
-

density
between the nuclei

lower energy than atomic orbital

Electrons in bonding molecular orbitals stabilize
a chemical bond.

MO Theory

Antibonding molecular orbital

electrons for the region between the nuclei

Highest electron density is located on
opposite sides of the nuclei

higher energy than atomic orbital

Electrons in antibonding molecular orbitals
destabilize a chemical bond.

MO Theory

Chemical Bonds

Chemical bond:

strong attractive force that
exists between atoms (or ions) in a compound

ionic bonds

covalent bonds

nonpolar covalent bond

polar covalent bond

Ionic Bond
:
the electrostatic force of
attraction between oppositely charged ions in
an ionic compound

metal cation (+)

non
-
metal anion (
-
)

Chemical Bonds

Covalent Bonds:

the attractive force between
atoms in a molecule that results from sharing
one or more pairs of electrons

non
-
metals

H
2
O, O
2
, CCl
4
, C
6
H
12
O
6

In some molecules, electrons are shared
equally.

nonpolar covalent bonds

H
-

H, Cl
-

Cl, O=O

Chemical Bonds

In some molecules, electrons are not shared
equally due to relatively “large” differences in
electronegativities between atoms in the bond.

polar covalent bonds

H
-

O

N
-

H

C
-

Cl

Electronegativity
:

tendency of an atom in a
compound to draw electrons towards itself

Chemical Bonds

Consider the C
-

Cl bond:

Cl is more electronegative than C

electrons are attracted more strongly to
Cl giving it higher electron density and a
partial negative charge
(
d
-

)

electrons are drawn away from C giving it
lower electron density and a
partial
positive charge
(
d
+

)

C Cl

en = 2.5 3.2

d
+

d
-

Polarity

The polarity of a bond is measured by its
dipole moment.

Amount of charge at either end of the
dipole x bond length

Common dipole moments:

C
-

H

0.3 D

C
-

O

0.86 D

N
-

H

1.31 D

C
-

Br

1.48 D

O
-

H

1.53 D

D = debye

Increasing
polarity

Chemical Bonds

How do you determine if a bond is polar?

As a rule of thumb:

D

en

Bond Type

< 0.5

nonpolar covalent

0.5
-

2.0

polar covalent

>
2.0

ionic

Note
:
These values are approximate. Bond length
is also important so there are exceptions to these
values!!!

Chemical Bonds

Polarity of bonds can be indicated in a couple
of ways:

partial charges (
d
+

and
d
-

)

d
+

on least electronegative atom

d
-

on most electronegative atom

a “+” sign at the positive end of the bond
and an arrow head at the negative end of
the bond

C Cl

Chemical Bonds

Example:

Which of the following contain polar
bonds? Identify all partial charges and indicate
the direction of the dipole moment for each polar
bond.

H
2
O, F
2
, HF, CH
3
CH
2
OH

Chemical Bonds

Example:

Which of the following contain polar
bonds? Identify all partial charges and indicate
the direction of the dipole moment for each polar
bond.

H
2
O
, F
2
,
HF
,
CH
3
CH
2
OH

d
-

d
+

d
+

d
-

d
+

Chemical Bonds

d
+

d
+

d
-

Chemical Bonds

Ionic and covalent compounds tend to have
different properties:

Ionic compounds tend to:

be water soluble

be solids at RT

have higher MP and higher BP

Covalent compounds tend to:

be less water soluble or completely
insoluble in water

be solids, liquids or gases at RT

have lower MP and lower BP

Chemical Bonds

Valence electrons

are involved in the formation
of chemical bonds or ions

electrons residing in the incomplete outer
shell of an atom

For
main group

elements, the number of
valence electrons for an element = group
number of the element

N (group 5A) has 5 valence electrons

Chemical Bonds

Lewis symbols

are used to depict the valence
electrons present in an atom (or ion).

chemical symbol for the element

dot for each valence electron

dots are placed on all 4 sides of the
chemical symbol

up to 2 dots (electrons) per side

Chemical Bonds

Example:

Draw the Lewis symbol for oxygen.

Chemical symbol:

O

Group number:

6A

# of valence electrons:

6

O