Chapter 1
Structure Determines Properties
Dr. Wolf's CHM 201 & 202
1- 1
Atoms, Electrons, and Orbitals
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1- 2
Atoms are composed of
Protons
+
positively charged
mass = 1.6726 X 10-27 kg
Neutrons
neutral
mass = 1.6750 X 10-27 kg
•
–
Electrons
negatively charged
mass = 9.1096 X 10-31 kg
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Atomic Number and Mass Number
A
Z
X
Atomic number (Z) = number of protons in nucleus
(this must also equal the number of electrons
in neutral atom)
Mass number (A) = sum of number of protons
+ neutrons in nucleus
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Schrödinger Equation
Schrödinger combined the idea that an electron
has wave properties with classical equations
of wave motion to give a wave equation for the
energy of an electron in an atom.
Wave equation (Schrödinger equation) gives a
series of solutions called wave functions ( ).
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Wave Functions
Only certain values of  are allowed.
Each  corresponds to a certain energy.
The probability of finding an electron at a
particular point with respect to the nucleus is
given by  2.
Each energy state corresponds to an orbital.
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Figure 1.1 Probability distribution ( 2) for an
electron in a 1s orbital.
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A boundary surface encloses the region
where the probability of finding an electron
is high—on the order of 90-95%
1s
2s
Figure 1.2 Boundary surfaces of a 1s orbital
and a 2s orbital.
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Quantum Numbers
Each orbital is characterized by a unique
set of four quantum numbers.
The principal quantum number (n) is a whole
number (integer, 1, 2, etc.) that specifies the shell
and is related to the energy of the orbital.
The angular momentum quantum number (l) is
usually designated by a letter (s, p, d, f, etc)
and describes the shape of the orbital. Values for l
range from 0 to (n-1)
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Quantum Numbers
Each orbital is characterized by a unique
set of four quantum numbers.
The magnetic quantum number m is a whole
number ranging from –l through 0 to +l. This gives
the orientation of the orbital.
The spin quantum number is either + or - ½
for the electron spin clockwise or counter clockwise
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s Orbitals
s Orbitals are spherically symmetric.
The energy of an s orbital increases with the
number of nodal surfaces it has.
A nodal surface is a region where the probability
of finding an electron is zero.
A 1s orbital has no nodes; a 2s orbital has one;
a 3s orbital has two, etc.
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The Pauli Exclusion Principle
No two electrons in the same atom can have
the same set of four quantum numbers.
Two electrons can occupy the same orbital
only when they have opposite spins.
There is a maximum of two electrons per orbital.
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First Period
Principal quantum number (n) = 1
Hydrogen
1s
Helium
Z=1
Z=2
1s 1
1s 2
2s
2p
H
He
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p Orbitals
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
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p Orbitals
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
There are three p orbitals for each value
of n (when n is greater than 1).
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p Orbitals
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
There are three p orbitals for each value
of n (when n is greater than 1).
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p Orbitals
p Orbitals are shaped like dumbells.
Are not possible for n = 1.
Are possible for n = 2 and higher.
There are three p orbitals for each value
of n (when n is greater than 1).
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Second Period
Principal quantum number (n) = 2
Z
1s
2s
2p
Li 3
Be 4
B 5
C 6
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Second Period
Z
N
7
O
8
F
9
1s
2s
2p
Ne 10
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Electron Configurations
Electrons fill the lowest energy levels first
(calcium shown)
4d
E
N
E
R
G
Y
Dr. Wolf's CHM 201 & 202
5s
4p
3d
4s
3s
2s
1s
3p
2p
Ionic Bonds
Positively charged ions called CATIONS
Negatively charged ions called XXX
DOGIONS
IONS
ANIONS
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Ionic Bonding
An ionic bond is the force of electrostatic
attraction between oppositely charged ions
Na+ (cation)
Cl– (anion)
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Ionic Bonding
Ionic bonds are common in inorganic chemistry
but rare in organic chemistry.
Carbon shows less of a tendency to form cations
than metals do, and less of a tendency to form
anions than nonmetals.
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Covalent Bonds and the Octet Rule
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The Lewis Model of Chemical Bonding
In 1916 G. N. Lewis proposed that atoms
combine in order to achieve a more stable
electron configuration.
Maximum stability results when an atom
is isoelectronic with a noble gas.
An electron pair that is shared between
two atoms constitutes a covalent bond.
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Covalent Bonding in H2
Two hydrogen atoms, each with 1 electron,
H.
.H
can share those electrons in a covalent bond.
H: H
Sharing the electron pair gives each hydrogen
an electron configuration analogous to helium.
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Covalent Bonding in F2
Two fluorine atoms, each with 7 valence electrons,
..
..
. F:
: ..F .
..
can share those electrons in a covalent bond.
.. ..
: ..
F : ..
F:
Sharing the electron pair gives each fluorine an
electron configuration analogous to neon.
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The Octet Rule
In forming compounds, atoms gain, lose, or
share electrons to give a stable electron
configuration characterized by 8 valence
electrons.
.. ..
: ..
F : ..
F:
The octet rule is the most useful in cases
involving covalent bonds to C, N, O, and F.
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Example
Combine carbon (4 valence electrons) and
four fluorines (7 valence electrons each)
.
. C.
.
..
: ..
F.
to write a Lewis structure for CF4.
..
.. : ..F: ..
:
:
F
: ..F: C
..
..
: ..F:
The octet rule is satisfied for carbon and
each fluorine.
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Example
It is common practice to represent a covalent
bond by a line. We can rewrite
..
.. : ..F: ..
:
: ..F: C
.. : ..F
: ..F:
..
: F:
as
..
: ..
F
C
..
..F:
: ..F:
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Double Bonds and Triple Bonds
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Inorganic Examples
..
..
: O: : C : : O:
..
:O
C
..
O:
C
N:
Carbon dioxide
H : C : :: N:
H
Hydrogen cyanide
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Organic Examples
H
.. H
..
H: C : : C:H
H
Ethylene
H
C
H
H
H : C : :: C:H
Dr. Wolf's CHM 201 & 202
Acetylene
H
C
C
C
H
1- 33
Polar Covalent Bonds
and Electronegativity
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Electronegativity
Electronegativity is a measure of the ability
of an element to attract electrons toward
itself when bonded to another element.
An electronegative element attracts electrons.
An electropositive element releases electrons.
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Pauling Electronegativity Scale
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
Electronegativity increases from left to right
in the periodic table.
Electronegativity decreases going down a group.
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Generalization
The greater the difference in electronegativity
between two bonded atoms; the more polar the
bond.
.. ..
: N N:
: ..
H—H
F ..
F:
nonpolar bonds connect atoms of
the same electronegativity
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Generalization
The greater the difference in electronegativity
between two bonded atoms; the more polar the
bond.

H
.. 
F:
..

H

..

O H
..
 
:O C
..

O
.. :
polar bonds connect atoms of
different electronegativity
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Electrostatic Potential Maps
Electrostatic potential maps show the charge
distribution within a molecule.

H
.. 
F:
..
Solid
surface
Red is negative charge;
blue is positive.
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Electrostatic Potential Maps
Electrostatic potential maps show the charge
distribution within a molecule.

H
.. 
F:
..
Transparent
surface
Red is negative charge;
blue is positive.
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Electrostatic Potential Maps
Electrostatic potential maps show the charge
distribution within a molecule.
.. 



H ..
H Li
F:
Red is negative charge;
blue is positive.
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Formal Charge
Formal charge is the charge calculated
for an atom in a Lewis structure on the
basis of an equal sharing of bonded
electron pairs.
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Nitric acid
Formal charge of H
H
..
O
..
..
O:
N
:O
.. :
We will calculate the formal charge for each
atom in this Lewis structure.
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Nitric acid
Formal charge of H
H
..
O
..
..
O:
N
:O
.. :
•Hydrogen shares 2 electrons with oxygen.
•Assign 1 electron to H and 1 to O.
•A neutral hydrogen atom has 1 electron.
•Therefore, the formal charge of H in nitric acid
is 0.
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Nitric acid
Formal charge of O
H
..
O
..
..
O:
N
: O:
..
•Oxygen has 4 electrons in covalent bonds.
•Assign 2 of these 4 electrons to O.
•Oxygen has 2 unshared pairs. Assign all 4 of
these electrons to O.
•Therefore, the total number of electrons
assigned to O is 2 + 4 = 6.
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Nitric acid
Formal charge of O
H
..
O
..
..
O:
N
:O
.. :
•Electron count of O is 6.
•A neutral oxygen has 6 electrons.
•Therefore, the formal charge of O is 0.
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Nitric acid
Formal charge of O
H
..
O
..
..
O:
N
:O
.. :
•Electron count of O is 6 (4 electrons from
unshared pairs + half of 4 bonded electrons).
•A neutral oxygen has 6 electrons.
•Therefore, the formal charge of O is 0.
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Nitric acid
Formal charge of O
H
..
O
..
..
O:
N
:O
.. :
•Electron count of O is 7 (6 electrons from
unshared pairs + half of 2 bonded electrons).
•A neutral oxygen has 6 electrons.
•Therefore, the formal charge of O is -1.
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Nitric acid
Formal charge of N
H
..
O
..
..
O:
N
–
:
:O
..
•Electron count of N is 4 (half of 8 electrons in
covalent bonds).
•A neutral nitrogen has 5 electrons.
•Therefore, the formal charge of N is +1.
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Nitric acid
Formal charges
H
..
O
..
..
O:
N+
–
:
:O
..
•A Lewis structure is not complete unless formal
charges (if any) are shown.
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Formal Charge
An arithmetic formula for calculating formal charge.
Formal charge =
group number
number of
number of
–
–
in periodic table
bonds
unshared electrons
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Formal Charge
"Electron counts" and formal
charges in NH4+ and BF41
H
+
H
4
Dr. Wolf's CHM 201 & 202
N
H
H
..
: F:
..
– ..
: ..
F B ..F:
: ..F:
7
4
1- 52
Structural Formulas of Organic Molecules
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Constitution
The order in which the atoms of a
molecule are connected is called its
constitution or connectivity.
The constitution of a molecule must be
determined in order to write a Lewis
structure.
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Condensed structural formulas
Lewis structures in which many (or all)
covalent bonds and electron pairs are
omitted.
H
H
H
H
C
C
C
H : O: H
H
Dr. Wolf's CHM 201 & 202
H
can be condensed to:
CH3CHCH3 or (CH3)2CHOH
OH
1- 55
Bond-line formulas
CH3CH2CH2CH3 is shown as
CH3CH2CH2CH2OH is shown as
OH
Omit atom symbols. Represent
structure by showing bonds between
carbons and atoms other than
hydrogen.
Atoms other than carbon and hydrogen
are called heteroatoms.
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Bond-line formulas
H
Cl
Cl
C
H2C
CH2
H2C
CH2
is shown as
C
H
H
Omit atom symbols. Represent
structure by showing bonds between
carbons and atoms other than
hydrogen.
Atoms other than carbon and hydrogen
are called heteroatoms.
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Constitutional Isomers
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Constitutional isomers
Isomers are different compounds that
have the same molecular formula.
Constitutional isomers are isomers
that differ in the order in which the
atoms are connected.
Another term for constitutional
isomers is “structural isomers.”
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A Historical Note
O
NH4OCN
Ammonium cyanate
H2NCNH2
Urea
In 1823 Friedrich Wöhler discovered that
when ammonium cyanate was dissolved in hot
water, it was converted to urea.
Ammonium cyanate and urea are
constitutional isomers of CH4N2O.
Ammonium cyanate is “inorganic.” Urea is
“organic.” Wöhler is credited with an important
early contribution that helped overturn the
theory of “vitalism.”
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Examples of constitutional isomers
H
H
C
H
..
O:
H
N+
:O
..
H
–
:
Nitromethane
C
..
O
..
N
..
..
O:
H
Methyl nitrite
Both have the molecular formula CH3NO2 but
the atoms are connected in a different order.
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Resonance
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Resonance
two or more Lewis structures may be
written for certain compounds (or ions)
Recall from Table 1.5
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Table 1.5 How to Write Lewis Structures
If an atom lacks an octet, use electron
pairs on an adjacent atom to form a
double or triple bond.
Example:
Nitrogen has only 6 electrons in the
structure shown.
H
..
..
:
H C O
N
O
..
..
..
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H
1- 64
Table 1.5 How to Write Lewis Structures
If an atom lacks an octet, use electron
pairs on an adjacent atom to form a
double or triple bond.
Example:
All the atoms have octets in this Lewis
structure.
H
..
..
:
H C O
N
O
..
..
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H
1- 65
Table 1.5 How to Write Lewis Structures
Calculate formal charges.
Example:
None of the atoms possess a formal
charge in this Lewis structure.
H
..
..
:
H C O
N
O
..
..
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H
1- 66
Table 1.5 How to Write Lewis Structures
Calculate formal charges.
Example:
This structure has formal charges; is
less stable Lewis structure.
H
.. –
+
:
H C O
N
O
..
..
..
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H
1- 67
Resonance Structures of Methyl Nitrite
same atomic positions
differ in electron positions
H
H
C
H
..
O
..
N
..
H
more stable
Lewis
structure
Dr. Wolf's CHM 201 & 202
..
O:
H
C
+
O
..
N
..
.. –
O
.. :
H
less stable
Lewis
structure
1- 68
Resonance Structures of Methyl Nitrite
same atomic positions
H
H
C
..
O
..
differ in electron positions
H
..
+
N
O:
H C O
..
..
H
more stable
Lewis
structure
Dr. Wolf's CHM 201 & 202
N
..
.. –
O
.. :
H
less stable
Lewis
structure
1- 69
Why Write Resonance Structures?
Electrons in molecules are often delocalized
between two or more atoms.
Electrons in a single Lewis structure are
assigned to specific atoms-a single Lewis structure
is insufficient to show electron delocalization.
Composite of resonance forms more accurately
depicts electron distribution.
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Example
Ozone (O3)
Lewis structure of
ozone shows one
double bond and
one single bond
••
•O
•
+
O
••
•• –
O ••
••
Expect: one short bond and one
long bond
Reality: bonds are of equal length
(128 pm)
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Example
Ozone (O3)
Lewis structure of
ozone shows one
double bond and
one single bond
••
•O
•
+
O
•• –
O ••
••
••
Resonance:
••
•O
•
Dr. Wolf's CHM 201 & 202
+
O
••
•• –
O ••
••
– ••
•O
• ••
+
O
••
O ••
••
1- 72
Example
Ozone (O3)
Electrostatic potential
map shows both end
carbons are equivalent
with respect to negative
charge. Middle atom
is positive.
••
•O
•
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+
O
••
•• –
O ••
••
– ••
•O
• ••
+
O
••
O ••
••
1- 73
The Shapes of Some Simple Molecules
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Valence Shell Electron Pair Repulsions
The most stable arrangement of groups
attached to a central atom is the one that has
the maximum separation of electron pairs
(bonded or nonbonded).
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Table 1.6 Methane
tetrahedral geometry
H—C—H angle = 109.5°
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Table 1.6 Methane
tetrahedral geometry
each H—C—H angle = 109.5°
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Table 1.6 Water
bent geometry
H—O—H angle = 105°
H
H
O
:
..
but notice the tetrahedral arrangement
of electron pairs
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Table 1.6 Ammonia
trigonal pyramidal geometry
H—N—H angle = 107°
H
H
N
:
H
but notice the tetrahedral arrangement
of electron pairs
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Table 1.6 Boron Trifluoride
F—B—F angle = 120°
trigonal planar geometry
allows for maximum separation
of three electron pairs
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Multiple Bonds
Four-electron double bonds and six-electron
triple bonds are considered to be similar to a
two-electron single bond in terms of their spatial
requirements.
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Table 1.6: Formaldehyde
H—C—H and H—C—O
angles are close to 120°
trigonal planar geometry
H
C
O
H
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Table 1.6 Carbon Dioxide
O—C—O angle = 180°
linear geometry
O
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C
O
1- 83
Molecular Dipole Moments
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Dipole Moment
A substance possesses a dipole moment
if its centers of positive and negative charge
do not coincide.
=exd
(expressed in Debye units)
+
—
not polar
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Dipole Moment
A substance possesses a dipole moment
if its centers of positive and negative charge
do not coincide.
=exd
(expressed in Debye units)
—
+
polar
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Molecular Dipole Moments
-
O
+
C
O
-
molecule must have polar bonds
necessary, but not sufficient
need to know molecular shape
because individual bond dipoles can cancel
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Molecular Dipole Moments
O
C
O
Carbon dioxide has no dipole moment;  = 0 D
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Figure 1.7
Carbon tetrachloride
=0D
Dr. Wolf's CHM 201 & 202
Dichloromethane
 = 1.62 D
1- 89
Figure 1.7
Resultant of these
two bond dipoles is
Resultant of these
two bond dipoles is
=0D
Carbon tetrachloride has no dipole
moment because all of the individual
bond dipoles cancel.
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Figure 1.7
Resultant of these
two bond dipoles is
Resultant of these
two bond dipoles is
 = 1.62 D
The individual bond dipoles do not
cancel in dichloromethane; it has
a dipole moment.
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Acids and Bases:
The Arrhenius View
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Definitions
Arrhenius
An acid ionizes in water to give protons. A
base ionizes in water to give hydroxide ions.
Brønsted-Lowry
An acid is a proton donor. A base is a proton
acceptor.
Lewis
An acid is an electron pair acceptor. A base
is an electron pair donor.
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Arrhenius Acids and Bases
An acid is a substance that ionizes to give
protons when dissolved in water.
H
A
.
H + + . A–
A base is a substance that ionizes to give
hydroxide ions when dissolved in water.
M
Dr. Wolf's CHM 201 & 202
..
OH
..
–. ..
M+ + . OH
..
1- 94
Arrhenius Acids and Bases
Strong acids dissociate completely in water.
Weak acids dissociate only partially.
H
A
.
H + + . A–
Strong bases dissociate completely in water.
Weak bases dissociate only partially.
M
Dr. Wolf's CHM 201 & 202
..
OH
..
–. ..
M+ + . OH
..
1- 95
Acid Strength is Measured by pKa
H
.
H + + . A–
A
[H+][A–]
Ka =
[HA]
pKa = – log10Ka
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Acids and Bases:
The Brønsted-Lowry View
Brønsted-Lowry definition
an acid is a proton donor
a base is a proton acceptor
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A Brønsted Acid-Base Reaction
A proton is transferred from the acid to the
base.
B .. + H A
base
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+
B
–
.
.
H + A
acid
1- 98
A Brønsted Acid-Base Reaction
A proton is transferred from the acid to the
base.
B .. + H A
base
Dr. Wolf's CHM 201 & 202
acid
+
B
–
.
.
H + A
conjugate
acid
conjugate
base
1- 99
Proton Transfer from HBr to Water
hydronium ion
H
..
..
.. O ..
Br
H
+
..
H
base
Dr. Wolf's CHM 201 & 202
acid
H
..–.
.. O+ H
.. Br
+
.. .
H
conjugate conjugate
acid
base
1- 100
Equilibrium Constant for Proton Transfer
H
..
.. O .. + H Br ..
..
H
Ka =
H
.. . –
.. O+ H + .. Br
.. .
H
[H3O+][Br–]
[HBr]
Takes the same form as for Arrhenius Ka, but
H3O+ replaces H+. H3O+ and H+ are
considered equivalent, and there is no
difference in Ka values for Arrhenius and
Brønsted acidity.
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1- 101
Equilibrium Constant for Proton Transfer
H
..
.. O .. + H Br ..
..
H
Ka =
H
.. . –
.. O+ H + .. Br
.. .
H
[H3O+][Br–]
[HBr]
pKa = – log10 Ka
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1- 102
Water as a Brønsted Acid
H
..
–
.. N ..
+ H OH
..
H
base
Dr. Wolf's CHM 201 & 202
acid
H
.. N
– ..
H + .. OH
..
H
conjugate conjugate
acid
base
1- 103
Dissociation Constants (pKa) of Acids*
stronger
acid
Acid
HI
weaker
acid
pKa
-10.4
Conj. Base
I–
HBr
-5.8
Br–
H2SO4
-4.8
HCl
H3O+
-3.9
HSO4–
Cl–
-1.7
H2O
strong acids are stronger than hydronium ion
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1- 104
Important Generalization!
stronger
acid
Acid
HI
-10.4
HBr
weaker
acid
pKa
-5.8
Conj. Base
–
I
–
Br
H2SO4
-4.8
HCl
-3.9
Cl
-1.7
H2O
+
H3O
HSO4
–
–
The stronger the acid, the weaker the conjugate base.
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1- 105
Dissociation Constants (pKa) of Acids*
Acid
pKa
Conj. Base
H3O+
–1.7
H2O
HF
3.5
F–
CH3CO2H
4.6
CH3CO2–
NH4+
9.2
NH3
H2O
15.7
HO–
weak acids are weaker than hydronium ion
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1- 106
Dissociation Constants (pKa) of Acids*
Acid
pKa
Conj. Base
CH3OH
15.2
CH3O–
H2O
15.7
HO–
CH3CH2OH
~16
CH3CH2O–
(CH3)2CHOH
~17
(CH3)3COH
~18
(CH3)2CHO–
(CH3)3CO–
alcohols resemble water in acidity; their conjugate
bases are comparable to hydroxide ion in basicity
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1- 107
Dissociation Constants (pKa) of Acids*
Acid
pKa
Conj. Base
NH3
~36
NH2–
(CH3)2NH
~36
(CH3)2N–
ammonia and amines are very weak acids;
their conjugate bases are very strong bases
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1- 108
Dissociation Constants (pKa) of Acids*
Acid
HC
H
pKa
CH
H
H
26
H2C
43
H
CH2
CH3CH3
HC
C
H
H
H
Conj. Base
H
–
H
H
45
62
–
H2C
H
–
CH
–
CH3CH2
Most hydrocarbons are extremely weak acids.
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1- 109
What Happened to pKb?
Dr. Wolf's CHM 201 & 202
1- 110
About pKa and pKb
A separate “basicity constant” Kb is not
necessary.
Because of the conjugate relationships in the
Brønsted-Lowry approach, we can examine
acid-base reactions by relying exclusively on
pKa values.
Dr. Wolf's CHM 201 & 202
1- 111
H
Example
H
N••
H
N
••
Which is the stronger base, ammonia (left) or
pyridine (right)?
Recall that the stronger the acid, the weaker the
conjugate base.
Therefore, the stronger base is the conjugate of
the weaker acid.
Look up the pKa values of the conjugate acids of
ammonia and pyridine in Table 1.7.
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1- 112
Example
H
H
+
N
H
pKa = 9.3
weaker acid
pKa = 5.2
stronger acid
H
+
N
H
Dr. Wolf's CHM 201 & 202
Therefore, ammonia is a
stronger base than pyridine
1- 113
How Structure Affects Acid Strength
Dr. Wolf's CHM 201 & 202
1- 114
The Main Ways Structure Affects Acid Strength
The strength of the bond to the atom from
which the proton is lost.
The electronegativity of the atom from which
the proton is lost.
Changes in electron delocalization on
ionization.
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1- 115
Bond Strength
Bond strength is controlling factor when
comparing acidity of hydrogen halides.
pKa
HF
HCl
HBr
HI
3.1
-3.9
-5.8
-10.4
weakest acid
strongest H—X bond
Dr. Wolf's CHM 201 & 202
strongest acid
weakest H—X bond
1- 116
Bond Strength
Recall that bond strength decreases in a group
in going down the periodic table.
Generalization: Bond strength is most
important factor when considering acidity of
protons bonded to atoms in same group of
periodic table (as in HF, HCl, HBr, and HI).
Another example: H2S (pKa = 7.0) is a
stronger acid than H2O (pKa = 15.7).
Dr. Wolf's CHM 201 & 202
1- 117
The Main Ways Structure Affects Acid Strength
The strength of the bond to the atom from
which the proton is lost.
The electronegativity of the atom from which
the proton is lost.
Changes in electron delocalization on
ionization.
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1- 118
Electronegativity
Electronegativity is controlling factor when
comparing acidity of protons bonded to atoms
in the same row of the periodic table.
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1- 119
Electronegativity
pKa
CH4
NH3
H2O
HF
60
36
15.7
3.1
weakest acid
least electronegative
Dr. Wolf's CHM 201 & 202
strongest acid
most electronegative
1- 120
Electronegativity
R
.. O .. + H A
R
.. O+ H + .. A–
H
H
The equilibrium becomes more favorable as A
becomes better able to bear a negative charge.
Another way of looking at it is that H becomes
more positive as the atom to which it is
attached becomes more electronegative.
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1- 121
Bond strength versus Electronegativity
Bond strength is more important when
comparing acids in which the proton that is lost
is bonded to atoms in the same group of the
periodic table.
Electronegativity is more important when
comparing acids in which the proton that is lost
is bonded to atoms in the same row of the
periodic table.
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1- 122
Acidity of Alcohols
In many acids
the acidic proton
is bonded to
oxygen.
Alcohols (RO—H)
resemble water
(HO—H) in their
acidity.
Dr. Wolf's CHM 201 & 202
pKa
HO—H
15.7
CH3O—H
15.2
CH3CH2O—H
16
(CH3)2CHO—H
17
(CH3)3CO—H
18
1- 123
Acidity of Alcohols
Electronegative substituents can increase the
acidity of alcohols by drawing electrons away
from the —OH group.
CH3CH2OH
CF3CH2OH
16
weaker
11.3
stronger
pKa
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1- 124
Inductive Effect
F
F
H
C
C
F
H
O
H
+
The greater acidity of CF3CH2OH compared to
CH3CH2OH is an example of an inductive effect.
Inductive effects arise by polarization of the
electron distribution in the bonds between
atoms.
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1- 125
Electrostatic Potential Maps
The greater positive character of the proton of
the OH group of CF3CH2OH compared to
CH3CH2OH is apparent in the more blue color
in its electrostatic potential map.
CH3CH2OH
Dr. Wolf's CHM 201 & 202
CF3CH2OH
1- 126
Another example of the inductive effect
O
CH3C O H
pKa
Dr. Wolf's CHM 201 & 202
4.7
weaker
O
CF3C
O H
0.50
stronger
1- 127
The Main Ways Structure Affects Acid Strength
The strength of the bond to the atom from
which the proton is lost.
The electronegativity of the atom from which
the proton is lost.
Changes in electron delocalization on
ionization.
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1- 128
Electron Delocalization
R
.. O .. + H A
R
.. O+ H + .. A–
H
H
Ionization becomes more favorable if electron
delocalization increases in going from right to
left in the equation.
Resonance is a convenient way to show
electron delocalization.
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1- 129
Nitric Acid
••
O ••
H
• O •• + H
•
••
O
••
N+
H
•• –
O ••
••
pKa = -1.4
••
O ••
H
•• O+ H +
H
Dr. Wolf's CHM 201 & 202
– ••
•O
•
••
N+
•• –
O ••
••
1- 130
Nitric Acid
••
Nitrate ion is stabilized by
electron delocalization.
Dr. Wolf's CHM 201 & 202
– ••
•O
•
••
O ••
N+
•• –
O ••
••
1- 131
Nitric Acid
•• –
•• O ••
–• ••
•O
••
Negative charge is
shared equally by all
three oxygens.
N+
O ••
••
•• –
•• O ••
•O
• ••
Dr. Wolf's CHM 201 & 202
N+
•• –
O ••
••
••
– ••
•O
•
••
O ••
N+
•• –
O ••
••
1- 132
Acetic Acid
••
O ••
H
• O •• + H
•
••
O
••
C
H
pKa = 4.7
CH3
••
O ••
H
•• O+ H +
H
Dr. Wolf's CHM 201 & 202
– ••
•O
•
••
C
CH3
1- 133
Acetic Acid
••
Acetate ion is stabilized by
electron delocalization.
– ••
•O
•
••
O ••
C
CH3
Dr. Wolf's CHM 201 & 202
1- 134
Acetic Acid
Negative charge is
shared equally by
both oxygens.
•• –
•• O ••
•O
• ••
C
– ••
•O
•
••
CH3
Dr. Wolf's CHM 201 & 202
••
O ••
C
CH3
1- 135
Acid-Base Equilibria
Dr. Wolf's CHM 201 & 202
1- 136
Generalization
The equilibrium in an acid-base reaction is
favorable if the stronger acid is on the left and
the weaker acid is on the right.
Stronger acid + Stronger base
Dr. Wolf's CHM 201 & 202
Weaker acid + Weaker base
1- 137
Example of a strong acid
H
..
.. O .. + H Br ..
..
H
pKa = -5.8
stronger acid
H
.. . –
.. O+ H + .. Br
.. .
H
pKa = -1.7
weaker acid
The equilibrium lies to the side of
the weaker acid. (To the right)
Dr. Wolf's CHM 201 & 202
1- 138
Example of a weak acid
••
O ••
H
••
•• O •• + H—OCCH
3
••
H
••
H
O ••
– ••
+
•• O—H + • OCCH
• ••
3
H
pKa = 4.7
weaker acid
pKa = -1.7
stronger acid
The equilibrium lies to the side of
the weaker acid. (To the left)
Dr. Wolf's CHM 201 & 202
1- 139
Important Points
A strong acid is one that is stronger than H3O+.
A weak acid is one that is weaker than H3O+.
A strong base is one that is stronger than HO–.
A weak base is one that is weaker than HO–.
The strongest acid present in significant
quantities when a strong acid is dissolved in
water is H3O+.
The strongest acid present in significant
quantities when a weak acid is dissolved in
water is the weak acid itself.
Dr. Wolf's CHM 201 & 202
1- 140
Predicting the Direction of Acid-Base Reactions
••
••–•
H—O • + H—OC
H
6
5
••
••
Phenol
pKa = 10
stronger acid
– ••
••
H—O—H + •• OC
H
6
5
••
••
Water
pKa = 15.7
weaker acid
The equilibrium lies to the side of the weaker
acid. (To the right) Phenol is converted to
phenoxide ion by reaction with NaOH.
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1- 141
Predicting the Direction of Acid-Base Reactions
O
••
••–•
HOCO• + H—OC
H
6
5
••
••
Phenol
pKa = 10
weaker acid
O
– ••
••
HOCO—H + •• OC
H
6
5
••
••
Carbonic acid
pKa = 6.4
stronger acid
The equilibrium lies to the side of the weaker
acid. (To the left) Phenol is not converted to
phenoxide ion by reaction with NaHCO3.
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1- 142
Lewis Acids and Lewis Bases
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1- 143
Definitions
Arrhenius
An acid ionizes in water to give protons. A
base ionizes in water to give hydroxide ions.
Brønsted-Lowry
An acid is a proton donor. A base is a proton
acceptor.
Lewis
An acid is an electron pair acceptor. A base
is an electron pair donor.
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1- 144
Lewis Acid-Lewis Base Reactions
The Lewis acid and the Lewis base can be
either a neutral molecule or an ion.
Lewis acid
+ Lewis base
A+
+
•• B–
A—B
A
+
•• B–
– A—B
A+
+
•B
•
A—B +
A
+
•• B
– A—B +
Dr. Wolf's CHM 201 & 202
1- 145
Example: Two Neutral Molecules
CH2CH3
F3B
+
•• O •
•
CH2CH3
Lewis acid
–
F3B
CH2CH3
+
O ••
CH2CH3
Lewis base
Product is a stable substance. It is a liquid with
a boiling point of 126°C. Of the two reactants,
BF3 is a gas and CH3CH2OCH2CH3 with a
boiling point of 34°C.
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1- 146
Example: Ion + Neutral molecule
•• –
H—O•• +
••
Lewis base
••
H3C—Br ••
••
••
H—OCH
3
••
+
•• –
•• Br•
•• •
Lewis acid
Reaction is classified as a substitution. But notice
how much it resembles a Brønsted acid-base reaction.
•• –
H—O•• +
••
Dr. Wolf's CHM 201 & 202
••
H—Br ••
••
••
H—O—H
••
+
•• –
• Br•
• •• •
1- 147
Example: Ion + Neutral molecule
•• –
H—O•• +
••
Lewis base
••
H3C—Br ••
••
••
H—OCH
3
••
+
•• –
•• Br•
•• •
Lewis acid
Brønsted acid-base reactions are a subcategory of
Lewis acid-Lewis base reactions.
•• –
H—O•• +
••
Dr. Wolf's CHM 201 & 202
••
H—Br ••
••
••
H—O—H
••
+
•• –
• Br•
• •• •
1- 148
End of Chapter 1
Dr. Wolf's CHM 201 & 202
1- 149