Chapter 1 Structure Determines Properties

1
Chapter 1Structure Determines Properties
2
1.1Atoms, Electrons, and Orbitals
3
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

4
Atomic Number and Mass Number

• 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

5
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
  aseries of solutions called wave functions (? ).

6
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
  isgiven by ? 2.
• Each energy state corresponds to an orbital.

7
Figure 1.1 Probability distribution (? 2) for an
electron in a 1s orbital.
8
A boundary surface encloses the regionwhere the
probability of finding an electronis highon the
order of 90-95.
Figure 1.2 Boundary surfaces of a 1s orbitaland
a 2s orbital.
9
Quantum Numbers

• Each orbital is characterized by a unique set
  of quantum numbers.
• The principal quantum number n is a wholenumber
  (integer) that specifies the shell and isrelated
  to the energy of the orbital.
• The angular momentum quantum number is usually
  designated by a letter (s, p, d, f, etc) and
  describes the shape of the orbital.

10
s Orbitals

• s Orbitals are spherically symmetric.
• The energy of an s orbital increases with
  thenumber of nodal surfaces it has.
• A nodal surface is a region where the
  probabilityof finding an electron is zero.
• A 1s orbital has no nodes a 2s orbital has
  onea 3s orbital has two, etc.

11
The Pauli Exclusion Principle

• No two electrons in the same atom can havethe
  same set of four quantum numbers.
• Two electrons can occupy the same orbitalonly
  when they have opposite spins.
• There is a maximum of two electrons per orbital.

12
First Period

• Principal quantum number (n) 1
• Hydrogen Helium
• Z 1 Z 2
• 1s 1 1s 2

H
He
13
p Orbitals

• p Orbitals are shaped like dumbells.
• Are not possible for n 1.
• Are possible for n 2 and higher.

14
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).

15
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).

16
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).

17
Second Period

• Principal quantum number (n) 2

18
Second Period
19
1.2Ionic Bonds
20
Ionic Bonding

• An ionic bond is the force of electrostaticattrac
  tion between oppositely charged ions.

21
Ionic Bonding

• Ionic bonds are common in inorganic
  chemistrybut rare in organic chemistry.
• Carbon shows less of a tendency to form
  cationsthan metals do, and less of a tendency to
  formanions than nonmetals.

22
1.3Covalent Bonds, Lewis Structures,and the
Octet Rule
23
The Lewis Model of Chemical Bonding

• In 1916 G. N. Lewis proposed that atomscombine
  in order to achieve a more stableelectron
  configuration.
• Maximum stability results when an atomis
  isoelectronic with a noble gas.
• An electron pair that is shared between two
  atoms constitutes a covalent bond.

24
Covalent Bonding in H2
Two hydrogen atoms, each with 1 electron,
can share those electrons in a covalent bond.

• Sharing the electron pair gives each hydrogen an
  electron configuration analogous to helium.

25
Covalent Bonding in F2
Two fluorine atoms, each with 7 valence electrons,
can share those electrons in a covalent bond.

• Sharing the electron pair gives each fluorine an
  electron configuration analogous to neon.

26
The Octet Rule
In forming compounds, atoms gain, lose, or share
electrons to give a stable electron configuration
characterized by 8 valence electrons.

• The octet rule is the most useful in cases
  involving covalent bonds to C, N, O, and F.

27
Example
Combine carbon (4 valence electrons) andfour
fluorines (7 valence electrons each)
to write a Lewis structure for CF4.
The octet rule is satisfied for carbon and each
fluorine.
28
Example
It is common practice to represent a
covalentbond by a line. We can rewrite
..
as
29
1.4Double Bonds and Triple Bonds
30
Inorganic Examples
Carbon dioxide
Hydrogen cyanide
31
Organic Examples
Ethylene
Acetylene
32
1.5Polar Covalent Bonds, Electronegativity,
and Bond Dipoles
33
Electronegativity
Electronegativity is a measure of the abilityof
an element to attract electrons toward itself
when bonded to another element.

• An electronegative element attracts electrons.
• An electropositive element releases electrons.

34
Pauling Electronegativity Scale
Electronegativity increases from left to right in
the periodic table. Electronegativity decreases
going down a group.
35
Generalization

• The greater the difference in electronegativityb
  etween two bonded atoms the more polar the
  bond.

HH
nonpolar bonds connect atoms ofthe same
electronegativity
36
Generalization

• The greater the difference in electronegativityb
  etween two bonded atoms the more polar the
  bond.

??
??
??


O
C
O
..
..
polar bonds connect atoms ofdifferent
electronegativity
37
Electrostatic Potential Maps

• Electrostatic potential maps show the
  chargedistribution within a molecule.

Solidsurface
Red is negative chargeblue is positive.
38
Electrostatic Potential Maps

• Electrostatic potential maps show the
  chargedistribution within a molecule.

Transparentsurface
Red is negative chargeblue is positive.
39
Electrostatic Potential Maps

• Electrostatic potential maps show the
  chargedistribution within a molecule.

?
?-
Li
H
Red is negative chargeblue is positive.
40
1.6Formal 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.

41
Nitric acid
Formal charge of H

• We will calculate the formal charge for each atom
  in this Lewis structure.

42
Nitric acid
Formal charge of H
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.
43
Nitric acid
Formal charge of 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.
44
Nitric acid
Formal charge of O

• Electron count of O is 6.
• A neutral oxygen has 6 electrons.
• Therefore, the formal charge of O is 0.

45
Nitric acid
Formal charge of 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.
46
Nitric acid
Formal charge of 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.
47
Nitric acid
Formal charge of N
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.
48
Nitric acid
Formal charges


..

• A Lewis structure is not complete unless formal
  charges (if any) are shown.

49
Formal Charge
An arithmetic formula for calculating formal
charge.
Group numberin periodic table
- Electron Count
Formal charge
number ofunshared electrons

Electron Count
½ (Number shared electrons)
50
Formal Charge
"Electron counts" and formal charges in NH4
and BF4-
7

4
51
1.7Structural Formulas of Organic Molecules
52
1.7Structural Formulas of Organic Molecules
53
1.7Structural Formulas of Organic Molecules
54
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.
55
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.
• Constitutional isomers are also called
  structural isomers.

56
Examples of constitutional isomers
Ethanol
Dimethyl ether

• Both have the molecular formula C2H6O but the
  atoms are connected in a different order.

57
Examples of constitutional isomers
..
H

O

H
N
C



O
H
..
Nitromethane
Methyl nitrite

• Both have the molecular formula CH3NO2 but the
  atoms are connected in a different order.

58
1.8Resonance
59
Resonance

• Two or more Lewis structures may be written for
  certain compounds (or ions).
• Recall from Table 1.6

60
Table 1.6 Introduction to the Rules of Resonance

• Step 1 The connectivity must be the same in
  all resonance structures.
• ExampleThe Lewis formulas below are not
  resonance forms of the same compound.

61
Table 1.6 Introduction to the Rules of Resonance

• Step 2 Each contributing structure must have
  the same number of electrons and same net charge.
• ExampleAll structures have 18 electrons and a
  net charge of 0.

62
Table 1.6 Introduction to the Rules of Resonance

• Step 3 Calculate formal charges on the first
  structure.
• ExampleNone of the atoms possess a formal
  charge in this Lewis structure.

63
Table 1.6 Introduction to the Rules of Resonance

• Step 4 Calculate formal charges on the second
  and third structures.
• ExampleThese structures have formal charges
  these are less stable Lewis structures.

64
Resonance Structures of Methyl Nitrite

• same atomic positions
• differ in electron positions

more stable Lewis structure
less stable Lewis structure
65
Resonance Structures of Methyl Nitrite

• same atomic positions
• differ in electron positions

more stable Lewis structure
less stable Lewis structure
66
Why Write Resonance Structures?

• Electrons in molecules are often
  delocalizedbetween 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.

67
Example

• Ozone (O3)
• Lewis structure of ozone shows one double bond
  and one single bond

Expect one short bond and one long
bond Reality bonds are of equal length (128 pm)
68
Example

• Ozone (O3)
• Lewis structure of ozone shows one double bond
  and one single bond

Resonance
69
Example

• Ozone (O3)
• Electrostatic potentialmap shows both
  endoxygens are equivalentwith respect to
  negativecharge. Middle oxygenis positive.

70
1.9Writing Organic Structures
71
Condensed structural formulas
Lewis structures in which many (or all) covalent
bonds and electron pairs are omitted.
can be condensed to
72
Bond-line formulas
Omit atom symbols. Represent structure by
showing bonds between carbons and atoms other
than hydrogen. Atoms other than carbon and
hydrogen are called heteroatoms.
73
Bond-line formulas
Cl
is shown as
simplifies to
74
1.10The Shapes of Some Simple Molecules
75
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).

76
Table 1.7 Methane

• tetrahedral geometry
• HCH angle 109.5

77
Table 1.7 Methane

• tetrahedral geometry
• each HCH angle 109.5

78
Table 1.7 Water

• bent geometry
• HOH angle 105

H
H

O
..
but notice the tetrahedral arrangement of
electron pairs
79
Table 1.7 Ammonia

• trigonal pyramidal geometry
• HNH angle 107

H
H

N
H
but notice the tetrahedral arrangement of
electron pairs
80
Table 1.7 Boron Trifluoride

• FBF angle 120
• trigonal planar geometry allows for maximum
  separationof three electron pairs

81
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
  spatialrequirements.

82
Table 1.7 Formaldehyde

• HCH and HCOangles are close to 120
• trigonal planar geometry.

83
Table 1.7 Carbon Dioxide

• OCO angle 180
• linear geometry

84
1.11Molecular Dipole Moments
85
Dipole Moment

• A substance possesses a dipole moment if its
  centers of positive and negative charge do not
  coincide.
• ? e x d
• (expressed in Debye units)

not polar
86
Dipole Moment

• A substance possesses a dipole moment if its
  centers of positive and negative charge do not
  coincide.
• ? e x d
• (expressed in Debye units)

polar
87
Molecular Dipole Moments
?
?-
?-

• molecule must have polar bonds
• necessary, but not sufficient
• need to know molecular shape
• because individual bond dipoles can cancel

88
Molecular Dipole Moments
Carbon dioxide has no dipole moment ? 0 D
89
Figure 1.7
Dichloromethane
Carbon tetrachloride
? 0 D
? 1.62 D
90
Figure 1.7
Resultant of thesetwo bond dipoles is
Resultant of thesetwo bond dipoles is
? 0 D
Carbon tetrachloride has no dipolemoment
because all of the individualbond dipoles cancel.
91
Figure 1.7
Resultant of thesetwo bond dipoles is
Resultant of thesetwo bond dipoles is
? 1.62 D
The individual bond dipoles do notcancel in
dichloromethane it hasa dipole moment.
92
1.12Curved Arrows and Chemical Reactions
93
Curved Arrows

• Curved arrows are used to track the flow of
• electrons in chemical reactions.
• The arrow begins where the electrons were
  originally and points to where they end up
• Consider the reaction shown below which shows
  the dissociation of A-B

94
Curved Arrows

• Consider the dissociation of H2CO3

95
Curved Arrows

• Many reactions involve both bond breaking
• and formation. More than one arrow may be
• required.

96
1.13Acids and BasesThe Arrhenius View
97
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.

98
Arrhenius Acids and Bases

• An acid is a substance that ionizes to give
  protons when dissolved in water.

A base is a substance that ionizes to give
hydroxide ions when dissolved in water.
99
Arrhenius Acids and Bases

• Strong acids dissociate completely in water.
  Weak acids dissociate only partially.

H

A
Strong bases dissociate completely in water.
Weak bases dissociate only partially.

M

100
Acid Strength is Measured by pKa

H
A
pKa log10Ka
101
1.14Acids and BasesThe Brønsted-Lowry View

• Brønsted-Lowry definitionan acid is a proton
  donora base is a proton acceptor

102
A Brønsted Acid-Base Reaction

• A proton is transferred from the acid to the base.

.
.
B

H
A
H
A
B

base
acid
103
A Brønsted Acid-Base Reaction

• A proton is transferred from the acid to the base.

.
.
B

H
A
H
A
B

base
acid
conjugate acid
conjugate base
104
Proton Transfer from HBr to Water
hydronium ion
H
H
..

..

.
.
.
.
O
H
Br
H
Br
O


..
..
H
H

• base acid conjugate conjugate acid base

105
Equilibrium Constant for Proton Transfer
H
H
..

.
.

O

H
Br
H
O
..
H
H
H3OBr
Ka
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.

106
Equilibrium Constant for Proton Transfer
H
H
..

.
.

O

H
Br
H
O
..
H
H
H3OBr
Ka
HBr
pKa log10 Ka
107
Water as a Brønsted Acid
H
H
..

..

.
.
N

H
OH
H
OH
N

..
..
H
H

• base acid conjugate conjugate acid base

108
Dissociation Constants (pKa) of Acids

• strong acids are stronger than hydronium ion

For a more detailed list click here for Table 1.8
109
Important Generalization!

• The stronger the acid, the weaker the conjugate
  base.

For a more detailed list click here for Table 1.8
110
Dissociation Constants (pKa) of Acids

• weak acids are weaker than hydronium ion

For a more detailed list click here for Table 1.8
111
Dissociation Constants (pKa) of Acids

• alcohols resemble water in acidity their
  conjugatebases are comparable to hydroxide ion
  in basicity

For a more detailed list click here for Table 1.8
112
Dissociation Constants (pKa) of Acids

• ammonia and amines are very weak acidstheir
  conjugate bases are very strong bases

For a more detailed list click here for Table 1.8
113
Dissociation Constants (pKa) of Acids
Acid
p
K
C
on
j
.

B
a
se
a
26
43
45
62
CH3CH3

• Most hydrocarbons are extremely weak acids.

For a more detailed list click here for Table 1.8
114
Tab. 1.8
115
Tab. 1.8(Contd.)
116
1.15What Happened to pKb?
117
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.

118
Example

• 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 1.8.

119
Example
H
H

weaker acid
N
H
pKa 9.3
H
pKa 5.2
stronger acid
Therefore, ammonia is a stronger base than
pyridine
120
1.16How Structure Affects Acid Strength
121
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
• Electron delocalization in the conjugate base

122
Bond Strength

• Bond strength is controlling factor when
  comparing acidity of hydrogen halides.

123
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).

124
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
• Electron delocalization in the conjugate base

125
Electronegativity

• Electronegativity is controlling factor when
  comparing acidity of protons bonded to atoms in
  the same row of the periodic table.

126
Electronegativity
127
Electronegativity

• 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.

128
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.

129
Acidity of Alcohols
130
Acidity of Alcohols

• Electronegative substituents can increase the
  acidity of alcohols by drawing electrons away
  from the OH group.

CH3CH2OH
CF3CH2OH
16
11.3
pKa
weaker
stronger
131
Inductive Effect

• 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.

132
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
CF3CH2OH
133
Another Example of the Inductive Effect
4.7
0.50
pKa
weaker
stronger
134
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
• Electron delocalization in the conjugate base

135
Electron Delocalization

• 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.

136
Nitric Acid

pKa -1.4
137
Nitric Acid

• Nitrate ion is stabilized by electron
  delocalization.

138
Nitric Acid

• Negative charge is shared equally by all three
  oxygens.

139
Acetic Acid

H
O


O




H
CH3
pKa 4.7
140
Acetic Acid

• Acetate ion is stabilized by electron
  delocalization.

141
Acetic Acid

• Negative charge is shared equally by both
  oxygens.

O



O
C


CH3
142
1.17Acid-Base Equilibria
143
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
Weaker acid Weaker base
144
Example of a Strong Acid
H

O
H


H
The equilibrium lies to the side of the weaker
acid. (To the right)
145
Example of a Weak acid


The equilibrium lies to the side of the weaker
acid. (To the left)
146
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.

147
Predicting the Direction of Acid-Base Reactions




HOC6H5
OC6H5


HOH



The equilibrium lies to the side of the weaker
acid. (To the right) Phenol is converted to
phenoxide ion by reaction with NaOH.
148
Predicting the Direction of Acid-Base Reactions
The equilibrium lies to the side of the weaker
acid. (To the left) Phenol is not converted to
phenoxide ion by reaction with NaHCO3.
149
1.18Lewis Acids and Lewis Bases
150
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.

151
Lewis Acid-Lewis Base Reactions

• The Lewis acid and the Lewis base can be either a
  neutral molecule or an ion.

152
Example Two Neutral Molecules
F3B

Lewis acid
Lewis base
Product is a stable substance. It is a liquid
witha boiling point of 126C. Of the two
reactants,BF3 is a gas and CH3CH2OCH2CH3 is a
liquid with a boiling point of 34C.
153
Example Ion Neutral Molecule

H3CBr

Lewis base
Lewis acid
Reaction is classified as a substitution. But
noticehow much it resembles a Brønsted acid-base
reaction.
154
Example Ion Neutral Molecule

H3CBr

Lewis base
Lewis acid
Brønsted acid-base reactions are a subcategory
ofLewis acid-Lewis base reactions.