Review

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29 Οκτ 2013 (πριν από 4 χρόνια και 2 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


Destructive interference leads exclusion of
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