Review
Atomic Structure,
Chemical Bonding and
Intro. To Molecular
Polarity
I. Atomic Structure
Atoms are primarily
composed of 3 sub atomic
particles.
Sub Atomic
Particle
Charge
Mass(amu)
proton (p)
+1
1
neutron(n)
0
1
electron(e-)
-1
0
An atom is neutral if
# e-’s = # p’s.
• If a neutral atom gains extra
electron(s) then it becomes a
negatively charged species
called an anion.
• If a neutral atom loses
electron(s) then it becomes a
positively charged species
called a cation.
An atom is completely
characterized by two
numbers;
the atomic #(Z) and the
atomic mass # (A).
1) Atomic # (Z) - the # of
protons in the nucleus responsible for identity of
the element.
2) Mass # (A)- the total # of
protons plus neutrons.
Representing Atoms of an Element
• An atom may be represented as its
Symbol preceded by its subscripted
atomic number, Z, and its superscripted
atomic mass number, A.
A
Z
Symbol
In the case of the element Carbon
12
C
6
Fig. 5-6, p.125
Z
A
Fig. 5-7, p.125
Arrangement of the subatomic
particles within the atom
 At the center is the nucleus which contains the
protons and neutrons.
 electrons may be thought of as traveling in
concentric shells or energy levels about the
nucleus.
 the energy of the shells increase as one
proceeds away from the nucleus.
There is a max. # of e-’s
that can be accommodated
in each shell.
Shell
Max. capacity
e-’s
1
2
2
8
3
18
4
32
Shell diagram for neutral
atom of Phosphorus (P)
15 p
16 n
Further development of
atomic model.
• Each shell is composed of 1
or more subshells.
• Each shell has as many
subshells as its own number.
– 1st shell has 1 subshell.
– 2nd shell has 2 subshells.
– 3rd shell has 3 subshells.
– 4th shell has 4 subshells.
There are only four different
kinds of subshells.
These subshells are labeled, in
order of increasing energy, by
the letters s, p, d & f.
Each subshell can accomodate a
different # of e-’s
Energy
Increases
subshell
e- capacity
s
p
d
f
2
6
10
14
Thus the total capacity
of shell is distributed
amongst its subshells.
8
Shell/subshell diagram for phosphorus
15 p
16 n
1s
2s 2p 3s 3p
The ground state electron configuration for phosphorus:
1s2 2s2 2p6 3s2 3p3
Atomic subshells in order of
increasing energy, filling
order.
4f _____
NOTE:
4d _____
th shell is
Although
the
4
4p _____
higher in energy than the 3rd
3d _____
shell, not all subshells of the
4s _____
4th shell are higher in
3p _____
energy than all subshells of
rd shell. In fact, the
the
3
3s _____
rd
highest
subshell
of
the
3
2p _____
shell (3d) is higher in energy
2s _____
than the lowest subshell of
1s _____
the 4th shell (4s)
Further development of
atomic model
Our latest model of the atom
identifies electrons as dots traveling
about the nucleus in concentric
subshells. The truth is that we can
never know the exact position of an
electron at any point in time. In
1926, however, Erwin Schrödinger (of
University of Zurich) developed a
theory known as Quantum mechanics in
which he worked out a mathematical
expression to describe the motion of
an electron in terms of its energy.
Further development of atomic
model
These mathematical expressions are
called wave equations since they are
based upon the concept that e-’s show
properties not only of particles but
also of electromagnetic waves. These
wave equations have a series of
solutions called wave functions which
allow us to predict the volume of space
around a nucleus in which there is a
high probability of finding a
particular e-. This volume of space in
which an electron is most likely to be
found is called an orbital.
Now, to fully develop our theory of
atomic structure we must understand
that the subshells
(s, p, d, f)
of our earlier atomic model consist
of orbitals that are not all
concentric in shape. Furthermore,
any one orbital can only accommodate
2 e-’s. Consequently, the number of
orbitals that comprise a subshell
can easily be calculated by simply
dividing the subshell capacity by 2.
Number of Orbitals in each Subshell
• Any s subshell has a capacity of 2 e-’s
– The number of orbitals that comprise any s
subshell is 1.
• Any p subshell has a capacity of 6 e-’s
– The number of orbitals that comprise any p
subshell is 3.
• Any d subshell has a capacity of 10 e-’s
– The number of orbitals that comprise any d
subshell is 5.
• Any f subshell has a capacity of 14 e-’s
– The number of orbitals that comprise any f
subshell is 7.
Orbitals (s p d + f)
• All orbitals of the same kind have
the same 3 dimensional shape but
different sizes. The size
increases with the energy level.
All s subshells consist of one s
orbital that is spherically
symmetrical about the nucleus. An
s orbital can accommodate 2 eThis accounts for the 2e- capacity
of the s subshell
s Orbitals
1s ORBITAL
2s ORBITAL
Each p subshell actually consists of a
set of three p orbitals of equal energy;
px py p z .
• Each of the three p orbitals
is dumbbell shaped and all
are oriented in space
perpendicular to one another.
• The max. capacity of each p
orbital is 2e-.
• This accounts for the total
capacity of the p subshell as
being 6 e-’s.
Each p subshell consists of a set of
three p orbitals of equal energy, px py pz
Shown together the three
p orbitals look like
this:
The d subshell actually
consists of a set of five d
orbitals of equal energy. Each
d orbital can hold a maximum of
2e-. This accounts for the
total capacity of the d
subshell as being 10 e-’s. We
will not be focusing on the d
orbitals therefore their shapes
and names need not be
memorized.
However, FYI……..
Electron Spins
• Electrons spin on their axis
• Physics tells us that any charged species
that spins, generates a magnetic moment.
That is to say, it acts like a tiny bar magnet
with a North and a South Pole.
• Furthermore, the “Right Hand Rule” tells
us that if we wrap the fingers of our right
hand around the spinning species, in the
direction of the spin, then our thumb will
be pointing to the magnetic north.
S
N
N
S
Represeanting Electrons
• Therefore, because of their magnetic
moments, we generally represent
electrons using a single barbed arrow.
The tip of the arrow points to the magnetic
north of the electron.
Atomic Orbitals in order of
Increasing Energy
3d__
3d__
3d__
ENERGY
4s___
3px__ 3py__ 3pz__
3s ___
2px__ 2py__ 2pz__
2s___
1s___
3d__
3d__
Ground - state electron
configurations
• This refers to the lowest
energy arrangement of e-’s
in orbitals about the
nucleus.
• To obtain this ground state electron configuration
electrons are assigned to
the orbitals of the previous
slide according to the three
Rules for Filling Orbitals
• Always fill the lowest energy orbitals first.
• The two electrons that occupy any orbital
must have opposite spins.
• When filling orbitals of equal energy (those
of the p,d,or f subshells) put one electron
in each orbital with their spins parallel until
all are half filled, then go back and pair
them.
Orbital Electron Configurations
• Write the orbital electron configuration for P
1s2 2s2 2px2 2py2 2pz2 3s2 3px1 3py1 3pz1
• Write the orbital electron configuration for O
1s2 2s2
2px2 2py1 2pz1
Using the periodic table
to write electron
configurations
The P.T. is arranged such that
each horizontal row (period)
represents the filling of
orbitals in their proper order.
More information from
the Periodic Table
• The term valence electron
refers to the # of e-’s in
the outermost energy level
or shell of an atom.
For all main group elements the #
of the column (family) of the
Periodic Table in which the symbol
for the element occurs = the # of
valence electrons.
Number of
Element
Valence e- s
Na
1
B
3
Cl
7
Lewis Structures of
Atoms
These are shorthand techniques
for emphasizing the outer shell
or valence e-’s of an atom by
representing an atom as its
symbol surrounded by its
valence e-’s, the e-’s in the
atoms outermost shell. Note
that the symbol of the element
represents the nucleus plus all
inner shell e-’s.
Write Lewis dot
structures for carbon,
hydrogen, oxygen,
nitrogen and chlorine.
carbon
hydrogen
C
nitrogen
H
chlorine
oxygen
N
O
Cl
Why do atoms react
together to form
compounds?
Atoms react with one another to
form compounds in an attempt to
achieve the e- configuration of
their nearest noble gas
neighbor (family 8). The
reason for this is that the econfiguration of the noble
gases represents an extremely
stable situation.
Noble (Inert) Gases
Fig. 5-6, p.125
There are two ways in which
atoms can bond together so
as to achieve the econfiguration of their
nearest noble gas neighbor.
1. They can loose or gain the
necessary e-’s and thereby
become ions and ultimately form
ionic bonds.
2. Two or more atoms can share e’s and form covalent bonds.
Ionic Bonds
These are formed when ions
anions/cations of opposite
charge come together.
Generally ionic compounds
are formed between metals
(left of step) and nonmetals
(right of step).
Consider the formation of the
ionic compound magnesium
bromide.
Magnesium (Mg
) could
achieve the e- config. of
Neon by loosing 2e- .

Mg 
 Mg
+2
+ 2e-
isoelectronic
with Ne
Bromine could achieve the econfig. of krypton by gaining
one e-.

 Br 


+ 1e-
 
 Br





Note  Kr 

:
Consequently one magnesium
combines with two bromine
atoms to form MgBr2.
Note all atoms in MgBr2
are isoelectronic with
their nearest noble gas
neighbor.
Mg+2 + Br-1 = MgBr2
Covalent Bond
• A covalent bond results from the sharing of
an electron pair between two atoms.
– Whenever two atoms share a pair of e-’s, it
is as if each member of the bonded pair of
atoms has gained an extra electron.
– As atoms bond together to become
isoelectronic with their nearest noble gas
neighbors, covalent bonds generally occur
when two or more nonmetallic elements
(right of step) bond together because the
nearest noble gas neighbors for these
elements lies ahead of them. Consequently,
they all need to gain electrons to become
isoelectronic with their nearest noble gas
neighbors.
How many hydrogen atoms bond to one carbon atom?
Can become
isoelectronic with
He by gaining 1e-
C
+
Can become
isoelectro
nic with
Ne by
gaining
4e-
4H
H
H C H
H
Lewis
Structure for
covalent
molecule of
CH4
Kekulé or
Lewis
structure for
covalently
bonded
molecule
H
H
C H
H
Rules for Creating Lewis Structures
for more Complicated Molecules
• Connect all atoms using single bonds
• Add the total # of Valence Electons
• Subtract 2e-’s from the total # of Valence e-’s for each
single bond drawn in first step
• Sprinkle any remaining e-’s so as to make all atoms
isoelectronic with their nearest noble gas neighbors.
This usually means 8 e-’s. For H its 2 e-’s.
• If there are insufficient e-’s to accomplish the previous
step, make one or more nonbonded e- pairs perform
double duty by forming multiple bonds.
Now let’s build the Ammonia Molecule
NH3
Connect all atoms using single bonds
Total # Valence Electrons = 8
H
Subtract 2e-’s for each single bond:
H
8 – (3 x 2) = 2 e-’s
Sprinkle remaining 2 e-’s so that all
atoms are isoelectronic with their
nearest noble gas neighbor
N H
H
N H
H
Let’s Try CO2
• Connect all w single bonds
• Total Valence e-’s = 16
• Subtract 2e-’s for each bond
O C
O
O C
O
O C
O
– 16 – (2 x 2) = 12
• Sprinkle remaining e-’s so
that all atoms have 8e-’s.
Peripheral atoms first.
• If the octet cannot be
satisfied for all, force
nonbonded pairs to perform
double duty
Now Let’s Try HCN
• Connect all w single bonds
H
• Total Valence e-’s = 10
• Subtract 2e-’s for each bond
C
N
– 10 – (2 x 2) = 6 e-’s
• Sprinkle remaining 6 e-’s so
that all atoms have 8e-’s.
H
Peripheral atoms first.
• If the octet cannot be satisfied
for all (except H), force
nonbonded pairs to perform H
double duty
C
N
C
N
Let’s Look at the Water Molecule
+
2H
O
H
O
H
In the water molecule each oxygen is isoelectronic with:
Neon
In the water molecule each hydrogen is isoelectronic
with:
Helium
Now Let’s Try the Amino Acid Alanine NH2CH(CH3)COOH
H
N
H
O
C
C
H H C H
• Connect all w single bonds, be
careful not to exeeed the
normal valences (combining
capacities) for all atoms
• Total Valence e-’s = 36
• Subtract 2e-’s for each bond
H
O
H2N
– 36 – (2 x 12) = 12 e-’s
• Sprinkle remaining 12 e-’s so
that all atoms have 8e-’s.
• If the octet cannot be satisfied
for all (except H), force
nonbonded pairs to perform
double duty
CH
C
OH
R
H2N
CH
R
COOH
O
H
Amino Acids
• There are 20 different Amino Acids.
Amino Acids differ from one another only
in the nature of the R side chain.
• Different R Side groups gives different
Amino Acids
R side chain
I
H2H— C —COOH
I
H
Examples of Amino Acids
H
I
H2N—C —COOH
I
H
glycine
CH3
I
H2N—C —COOH
I
H
alanine
Different Types of R groups –
Different Amino Acids
Nonpolar R = H, CH3, alkyl groups, aromatic
O
Polar
ll
R = –CH2OH, –CH2SH, –CH2C–NH2,
(polar groups with –O-, -SH, -N-)
Polar/Acidic
R = –CH2COOH, or -COOH
Polar/ Basic
R = –CH2CH2NH2
Amino Acids and Proteins
• Amino Acids are the building blocks of
proteins
– In fact proteins are simply combinations of amino
acids linked in a head to tail fashon
Types of Proteins
•
•
•
•
•
•
•
Type
Structural
Contractile
Transport
Storage
Hormonal
Enzyme
Protection
Examples
tendons, cartilage, hair, nails
muscles
hemoglobin
milk
insulin, growth hormone
catalyzes reactions in cells
immune response
Kekulé or Lewis structure
for water molecule.
O
H
H
The Covalent Bond and
Electronegativity
• The sharing of an e- pair
between two atoms may be equal .
– If this is the case then the
resulting covalent bond is a
nonpolar covalent bond.
• If, on the other hand the
sharing is unequal then a polar
covalent bond results.
The reason for this
variance in bond polarity
is due to the fact that
different elements have
different tendencies to
attract to themselves extra
electrons. In other words,
each element has a
different electronegativity
Electronegativity
The tendency of an atom,
when in combination with
other atoms, to attract to
itself the bonded (extra) e’s.
Electronegativity values
increase from left to right
across any horizontal row
(period) of the P.T. and they
decrease going down any
vertical column (family) of
the P.T.
Consequently the most
electronegative elements are
N, O, F, Cl, Br
Electronegativity values for
selected elements.
• If two atoms are covalently bonded
and one has a high
electronegativity and the other has
a low electronegative then the
electron pair comprising that bond
is not shared equally but spends
more of its time closer to the more
electronegative atom. The
immediate result of this unequal
sharing is that the more
electronegative atom gains a
partial negative charge (-) while
the less electronegative element
gains a partial positive charge (
+).
This type of bond is called a
polar covalent bond.
The degree to which a covalent bond
is polarized is indicated by the
electronegativity difference
between the two bonded atoms.
Refer to next slide for
electronegativity values of
elements.
• If the electronegativity difference
is greater than .5 but less than 2.0
then the covalent bond is polar.
• If the electronegativity difference
is less than .5 then the covalent
bond is nonpolar.
Polar Covalent Bonds in H2O
Electronegativity Difference Between Oxygen and
Hydrogen is:
-
1.4
O
H
+
H +
A molecule typical of those found
in petroleum. The bonds are not
polar.
Electronegativity Difference Between Carbon and Hydrogen
is:
0.4
Ionic Bond and
Electronegativity
• Consideration of
electronegativity can
demonstrate that ionic bonds
are nothing more than an
extreme case of a polar
covalent bond. In fact…
• if the electronegativity
difference between two atoms
is greater than 2.0, then any
bond between these two atoms
would be ionic.
Molecular Polarity
If a molecule contains polar
bonds, and if those polar bonds
are located such that the  +
charges are at one end of the
molecule and the  - charges
are at the other end, then the
molecule is a polar molecule.
The measure of molecular
polarity is a quantity called
the dipole moment (D).
Like Dissolves Like
• Polar molecules dissolve in Polar Solvents
• Nonpolar molecules dissolve in nonpolar
solvents
• Polar molecules do not dissolve in
nonpolar solvent
• Nonpolar molecules do not dissolve in
polar solvents
: An oil
layer
floating
on water.
The oil is
nonpolar
and the
water is
polar
Polar water molecules interact with the positive and
negative ions of a salt. Ionicly bonded materials are
the extreme case of polar substances
CHAPTER 01 (FUNDAMENTALS
Org. Chem) CONTINUED
The Covalent Bond In Organic
Chemistry
• The covalent bond is of chief importance
in organic chemistry
The Covalent Bond and Valence
Bond Theory
• Valence Bond theory offers a description
of the covalent bond in terms of the orbital
model of the atom. Valence Bond theory
maintains that covalent bonds are formed
by an overlapping of two half-filled (1e-)
atomic orbitals.
• Consider the formation of the H2 molecule
from two isolated hydrogen atoms:
Formation of the H2 molecule
H
+
H
+
H
H
H
H
H H
The reaction is accompanied by the evolution of 104
kcal/mole H2 formed. This means that the product
(H2 molecule) is more stable than the reactants
(isolated H atoms) by 104 K cal/mole. The bond
strength of the H2 molecule is 104 kcal/mole. This
means that it would take 104 kcal to rupture the
bonds in 1 mole of H2 molecules.
The Valence Bond
representation of covalent bond
formation
1
Valence Bond Theory identifies two
types of Covalent Bonds
• Sigma (δ) Bonds – the bond in the H2
molecule is a sigma bond. Sigma bonds
result from the head to head overlap of
two half filled atomic orbitals. Sigma
bonds are cylindrically symmetrical about
a line joining the two nuclei
• Pi (π) Bonds – These result from the sideto-side overlap of two half filled atomic
orbitals.
Orbital overlap to form Sigma
(δ) Bonds
The Sigma Bond in the H2, HCl and the Cl2
molecules
Nucleus
Formation of the Pi (π) Bonds
Pi bonds are always accompanied by a Sigma Bond. A Pi
Bond cannot form without first forming a Sigma Bond.
Pi Bond
Consider the Formation of the O2
molecule
The Orbital electron configuration for oxygen is:
1s2
2s2 2px2 2py1 2pz1
+
O
+
O
O
O
Hybridization of Atomic Orbitals
• Certain atoms
their atomic
orbitals before bonding to other atoms and
forming molecules.
• The reason for this is that Hybridized Atomic
Orbitals are more directional and offer more
effective overlap than do unhybridized atomic
orbitals. As the strength of a covalent bond is
directly related to the extent of overlap of the
two ½ filled atomic orbitals, hybridized atomic
orbitals form stronger bonds than do
unhybridized atomic orbitals.
Evidence for the Hybridization of
Atomic Orbitals
• Consider the H2O molecule.
O
H
H
• We know that the electron pair geomery is:
Tetrahedral
• We therefore know that it’s HOH bond angle is:
109.5 degrees
However
• If a covalent bond results from the overlap of
two ½ filled atomic orbitals and if the orbital
electron configuration for Oxygen is:
1s2
2s2 2px2 2py1 2pz1
• Then the HOH bond angle should be 90
degrees and look like this:
H
H
or
H
O
H
NH3
and CH4
• The same proof of Hybridization can be
obtained by comparing the actual shapes and
bond angles in ammonia and methane to what
they would be if N and C used their
unhybridized atomic orbitals to bond the H’s.
• The orbital electron configurations for N and C
are: N; 1s2 2s2 2px1 2py12pz1 and C; 1s2 2s2
2px12py1
• In each case the expected bond angle would be
90 degrees as compared to the actual bond
angle of 109.5 degrees
Hybridizations States of Carbon
• Carbon can adopt any one of three hybridization
states, sp3( tetrahedral molecular geometry),
sp2(trigonal planar), sp (linear) depending upon the
number of electron pairs about the carbon.
• If C has 4 electron pairs as in CH4 then it sp3
hybridizes. Bond angles = 109.5°
• If C has three electron pairs as in ethene,
it sp2 hybridizes.
H
HBond angles= 120°
•
C
C
H
H
• If C has two electron pairs as in ethyne,
it sp hybridizes. 180°
H C C H
Practice Problems
• Identify the hybridization states and the
bond angles for each carbon atom in the
following molecule.
CH3CH2CH
sp3
sp3 sp2
C
sp
CHCH2C
sp2 sp3
All sp3 C’s 109.5 degrees
All sp2 C’s 120 degrees
All sp C’s 180 degrees
sp
CH
sp
sp3, sp2 and sp Hybridized Orbitals
• How are they formed from the atomic
orbitals?
• What do they look like?
The orbital electron configurations for C is: 1s2 2s2 2px1 2py1 2pz0
Formation of
four sp3 hybrid
orbitals
H
or
H
C
H
H
sp3 Hybrid Orbitals
• sp3 Hybrid orbitals are formed from one s and three p
orbitals. Therefore, there are four large lobes.
• Each lobe points towards the vertex of a tetrahedron.
• The angle between the large lobs is 109.5.
• All molecules with tetrahedral electron pair geometries
are sp3 hybridized.
sp3 Hybrid Orbitals
• Tetrahedral e- pair geometry
• 109.5° bond angle
=>
sp2 Hybrid Orbitals
• Important: when we mix n atomic orbitals we must get n
hybrid orbitals.
• sp2 hybrid orbitals are formed with one s and two p
orbitals. (Therefore, there is one unhybridized p orbital
remaining.)
• The large lobes of sp2 hybrids lie in a trigonal plane.
• All molecules with trigonal planar electron pair
geometries have sp2 orbitals on the central atom.
The orbital electron configurations for C is: 1s2 2s2 2px1 2py1 2pz0
These three atomic orbitals mix to form three sp2
hybrid orbitals
All together the 3 sp2’s and the unhybridized p look like this:
Formation of
three sp2 hybrid
orbitals
• Show Brown and LeMay Clip here
• The Shape of Molecule Chapter
sp Hybrid Orbitals
•sp hybrid orbitals are formed with one s and one p
orbitals. (Therefore, there are two unhybridized p orbital
remaining.)
The lobes of sp hybrid orbitals are 180º apart.
The orbital electron configurations for C is: 1s2 2s2 2px1 2py1 2pz0
These two atomic orbitals mix to form two sp hybrid
orbitals
Formation of
two sp hybrid
orbitals
All together the 2 sp’s
and the 2 unhybridized
p’s look like this:
Multiple Bonds
• A double bond (2 pairs of shared electrons)
consists of a sigma bond and a pi bond.
• A triple bond (3 pairs of shared electrons)
consists of a sigma bond and two pi bonds.
=>
Sample Problems
• Predict the hybridization, geometry, and
bond angle for each atom in the
following molecules:
• Caution! You must start with a good
Lewis structure!
• NH2NH2
• CH3-CC-CHO
O
CH3
C
_
CH2
=>
Acids and Bases: The Brønsted–
Lowry Definition
• The terms “acid” and “base” can have different
meanings in different contexts
• The Arrhenius definition of an acid is any substance
that increases the H+ conc, of water. An Arrhenius
base increases the OH- conc. of water.
• The idea that acids are aqueous solutions containing a
lot of “H+” and bases are solutions containing a lot of
“OH-” is not very useful in organic chemistry
• Instead, Brønsted–Lowry theory defines acids and
bases by their role in reactions that transfer protons
(H+) between donors and acceptors
Brønsted Acids and Bases
• A Brønsted acid is a substance that
donates a hydrogen ion (H+) also called a
proton.
• A Brønsted base is a substance that
accepts the H+
– “proton” is a synonym for H+ - loss of an
electron from H leaving the bare nucleus—a
proton
The Reaction of HCl with H2O
• When HCl gas dissolves in water, a Brønsted acid–base
reaction occurs
• HCl donates a proton to water molecule, yielding
hydronium ion (H3O+) and Cl
• The reverse rxn. is also a Brønsted acid–base reaction
between the conjugate acid and conjugate base
Acids are shown in red, bases in blue. Curved arrows go from
bases to acids
Conjugate Acids and Conjugate
Bases
• Every acid has association with it a
conjugate base formed from the acid by
loss of a proton.
– The conjugate base of HCl is Cl-.
• Every base has associated with it a
conjugate acid formed from the base by
addition of a proton.
– The conjugate acid of H2O is H3O+.
Acid Strengths
• Acid strengths are indicated by the extent to which
they donate protons to water. This extent of proton
donation is indicated by the equilibrium constant for
the reaction of the acid with water.
–
–
–
HA + H2O
Keg = [A-}{H3O+]
[HA}{H2O}
A- + H3O+
products
reactants
• The higher the keg, the greater the tendency for the
acid to donate protons to water. Therefore, the
higher the keg, the stronger the acid, the lower the
keg, the weaker the acid
Ka Values
• As the [H2O] remains constant for most keq
measurements, we may therefore rewrite
the equilibrium expression using a new
term called the acidity constant Ka.
•
Ka = keq [H2O] = [H3O+] [A-]
•
[HA]
• Therefore, the stronger the acid, the
greater the Ka, the weaker the acid, the
lower the Ka.
pKa Values
• Acid strengths are usually defined in terms
of pKa values
–
pKa = -log Ka
• The stronger the acid, the lower the pKa
and the weaker the acid, the higher the
pKa value.
Predicting Acid–Base Reactions
from pKa Values
• An Acid will react with a Base if and only if the
conjugate acid that is formed is weaker (has a higher
pKa value) than the original acid.
• This simple concept will allow one to predict whether
or not an acid/base reaction will go by simply
comparing the pKa values of the original acid and its
conjugate acid. Will the following reaction go?
Yes!
Predicting Acid Base Reactions
• Will the following reaction go?
O
CH3
C OH
pKa 4.74
O
+ CH3
NH2
pKb 3.36
Yes!
CH3
C O
pKb 9.26
-
+
CH3
+
NH3
pKa 10.64
Organic Acids and Organic Bases
• The reaction patterns of organic
compounds often are acid-base
combinations
• The transfer of a proton from a strong
Brønsted acid to a Brønsted base, for
example, is a very fast process and will
always occur along with other reactions
Organic Acids
• Those that lose a proton from O–H, such as
methanol and acetic acid
• Those that lose a proton from C–H, usually from
a carbon atom next to a C=O double bond
(O=C–C–H)
Organic Bases
• Have an atom with a lone pair of electrons that
can bond to H+
• Nitrogen-containing compounds derived from
ammonia are the most common organic bases
• Oxygen-containing compounds can react as
bases when with a strong acid or as acids with
strong bases
Acids and Bases: The Lewis
Definition
• Lewis acids are electron pair acceptors and
Lewis bases are electron pair donors
• The Lewis definition leads to a general
description of many reaction patterns but there is
no scale of strengths as in the Brønsted
definition of pKa
Illustration of Curved Arrows in Following
Lewis Acid-Base Reactions
Lewis Acids and the Curved Arrow
Formalism
• The Lewis definition of acidity includes metal cations,
such as Mg2+
– They accept a pair of electrons when they form a bond to a
base
• Group 3A elements, such as BF3 and AlCl3, are Lewis
acids because they have unfilled valence orbitals and
can accept electron pairs from Lewis bases
• Transition-metal compounds, such as TiCl4, FeCl3,
ZnCl2, and SnCl4, are Lewis acids
• Organic compounds that undergo addition reactions with
Lewis bases (discussed later) are called electrophiles
and therefore Lewis Acids
• The combination of a Lewis acid and a Lewis base can
shown with a curved arrow from base to acid
Lewis Acids and Bases
• Acids accept electron pairs = electrophile
• Bases donate electron pairs = nucleophile
CH2 CH2
nucleophile
+
BF3
electrophile
_
BF3
CH2
+
CH2
=>
Lewis Bases
• Most oxygen- and nitrogen-containing organic
compounds are Lewis bases because they have lone
pairs of electrons
• Some compounds can act as both acids and bases,
depending on the reaction
Drawing Chemical Structures
• Chemists use shorthand ways for writing structures
• Condensed structures: C-H and C-C and single bonds aren't
shown but understood
– If C has 3 H’s bonded to it, write CH3
– If C has 2 H’s bonded to it, write CH2; and so on. The compound called
2-methylbutane, for example, is written as follows:
• Horizontal bonds between carbons aren't shown in condensed
structures—the CH3, CH2, and CH units are simply but vertical
bonds are added for clarity
Skeletal Structures
• C’s are not shown. They are assumed to be at
each intersection of any two lines (bonds) and at
end of each line
• H’s bonded to C’s aren't shown –Since carbon
always has a valence of 4, we mentally supply
the correct number of H’s by subtracting the # of
bonds shown from 4.
• All atoms other than C and H are shown
• See next slide for examples
0
H
0
Multiple Bonds
In condensed formulas double and triple bonds are drawn as they
would be in a Lewis structure showing two dashes for a double
bond and three dashes for a triple bond