Fall 2004 Supplemental notes
Acids and Bases
“Curved Arrow Formalism” or Pushing Electrons
Carbon and other second row elements such as B, N, O, and F follow the octet
rule, i.e. they try to have the sum of bonding electrons and electrons in lone
pairs around them equal to 8. For the first row, hydrogen tries to have 2
electrons.
In general, NONE of these elements will have more than an octet (or duet for
hydrogen).
Electron Deficient Compounds
Sometimes molecules have atoms that are short of an octet by one or more
electron pairs – they tend to be very reactive. For example:
1. H + has 0 electrons and it needs 2, thus it is deficient by 2.
2. BF3 is an electron deficient compound. The boron atom in boron tri-fluoride
has 6 electrons, and it needs 8. Thus it is deficient by 2 electrons. One
additional lone pair is needed to fill its octet.
F
F
B
F
F
F
F
B
B
F
F
F
3. Methyl cation has 6 electrons, and it needs 8, thus is deficient by 2.
H
H
C
H
H
H
H
C
-1-
H
C
H
H
Fall 2004 Supplemental notes
Lewis Acids and Lewis Bases
F
F
F
F
B
+
F
B
F
F
B
F
F
F
Lewis Base
F
F
8 Electron
Tetrafluoroborate ion
Lewis Acid
Electron deficient compounds, which can behave a electron pair acceptors are
Lewis acids.
A species that donates an electron pair is a Lewis base.
The reaction above is called Lewis acid/ Lewis base association reaction.
Lewis acid ⇒ electrophile (“loves electrons“)
Lewis base ⇒ nucleophile (why??)
We will see many Lewis acid-Lewis base reactions in coming months.
IT IS VERY IMPORTANT to be able to identify Lewis acids and Lewis bases.
Another example (simplified):
lone
pair
H
N
H
-
H
H
8 e Lewis Base
+
N
H
H
-
0 e Lewis Acid
-2-
H
H
ammonium ion
8 e-
Fall 2004 Supplemental notes
What Do Curved Arrows Mean?
source
of electrons
destination
of electrons
F
BF3
F
new bond, electrons shared
BF3
The curved arrow indicates the flow of electrons.
The arrow always starts at the electron donor and ends at the electron acceptor.
Here the arrow starts at the Lewis base end and ends at the electron deficient
species (the Lewis acid).
**Note that charge is conserved.
What about the reverse reaction?
F
F
B
F
F
F
+
B
F
-3-
F
F
Fall 2004 Supplemental notes
Conjugate Acid-Base Pairs
H
N
H
Base
H
H
+
N
Br
H
H
H
H
Conjugate
acid
Acid
+
Br
Conjugate
base
Conjugate Acid-Base Pairs
1. Note that the H-Br bond is broken and NH3 -H bond is formed.
2. Why is the H-Br bond broken? Because the H in HBr already had a
duet and if it is to accept two electrons from ammonia, it must also lose
two.
3. When a lone pair is contributed, the formal charge on the atom
contributing the lone pair becomes more positive by one integer, and
when a lone pair is gained, the formal charge on the atom receiving the
lone pair becomes more negative by one integer.
Example:
0
N
H
H
H
H
0
+
H
N
Br
H
-1
+1
H
H
+
Br
Nitrogen contributes a lone pair to form a new bond, so the charge increases by 1.
Bromine gains a lone pair when the bond is broken, so the charge decreases by one.
Note that the net charge on both side of the arrow should be the same (charge is
conserved).
-4-
Fall 2004 Supplemental notes
Examples
Drawn to show tetrahedral
geometry with lone pairs
occupying sites.
H
H
O
CH3
CH3
O
Base
O
H
H
Conjugate
acid
Acid
+
O
CH3
CH3
Conjugate
base
Conjugate Acid-Base Pairs
• Note the charges, bonds formed and bonds broken.
• Note the conjugate Lewis acid pair and Lewis base pair.
• Note that the arrows indicate flow of electrons.
What about:
H
H
O
Base
CH3
CH3
O
O
H
H
Conjugate
acid
Acid
+
O
CH3
CH3
Conjugate
base
Conjugate Acid-Base Pairs
Wrong
Reaction is not wrong, BUT use of the curved arrow is incorrect.
REMEMBER: Electrons flow from tail to head!!
Also note BrØnsted-Lowry Acid and Bases:
• BrØnsted Acid – A species which reacts by donating a proton (H+).
• BrØnsted Base – A species that can accept a proton.
BrØnsted-Lowry summary:
m
H X
Bronsted
Acid
n
+
m-1
B
X
Bronsted
Base
+
Conjugate
Base
H
n+1
B
Conjugate
Acid
So BrØnsted-Lowry Acid-Base definition is a more limited definition than
Lewis acid base.
-5-
Fall 2004 Supplemental notes
Use of Curved Arrow Formalism to Derive Resonance Structures
Not all molecules can be described well by one Lewis structure.
In many cases another structure can be derived by a shift of one or more
electron pairs.
Both structures for benzene are equal in energy.
Which structure is correct?
Actually, neither is correct. The real structure of benzene is in between the two
structures above. The two structures shown above are called two limiting
resonance structures.
** Extremely important: Resonance does NOT imply rapid interchange
between structures, but rather that the actual structure is a weighted average of
the two (or more) limiting resonance structures.
Curved arrows can help one draw resonance structures.
Here the arrow describes ‘flow’ in the loose sense of the word.
circle impies 1.5 bonds
between carbons
-6-
Fall 2004 Supplemental notes
Other Examples
O
O
H3C
H3C
O
O
O
H3C
O
curve implies 1.5 bonds
between carbons
Note, like in the case of Lewis acids-Lewis bases:
•
•
•
•
•
The arrow represents “flow” of electron pair.
Flow "in", means formation of new bond;
Flow "out", means breaking a bond.
Atoms should not violate octet rule.
The overall charge is conserved.
O
O
H3C
H3C
N
CH3
N
CH3
CH3
CH3
Left structure, no charge separation.
Right structure +,- so right structure is higher energy and contributes less.
O
H2C
O
CH3
H2C
CH3
Right structure has the minus charge on more electronegative atom.
So, the right structure is lower in energy and contributes more to the actual
structure of the molecule.
-7-
Fall 2004 Supplemental notes
Stability of Resonance Structures and Summary for Deriving Them:
•
•
•
•
•
•
•
•
Try to satisfy octet.
Maximize the number of covalent bonds.
Minimize charge separation.
Try to place negative charges (electrons) on most electronegative
atom.
Positive charge on halogens is really bad (because they are highly
electronegative).
Fewer than four bonds to carbon is quite bad.
Charges on carbon are quite bad.
More than 8 electrons on carbon, nitrogen, or oxygen, is
unacceptable.
-8-
Fall 2004 Supplemental notes
BrØnsted -Lowry Acid Base Equilibria
Equilibrium constants:
X
H
n
+
B
m
n-1
X
+
H
B
m+1
e.g.
Cl
H
+
OH
Cl
+
H
OH
One can write an equilibrium expression:
Keq =
X
n-1
H
n
X
H
B
m+1
Molarity of species
m
B
Keq > 1 implies reaction goes to the right
Keq < 1 implies reaction goes to the left
Keq > 1 implies that X-Hn is a stronger acid than H-Bm+1
and that :Bm is a stronger base than Xn-1
-9-
Fall 2004 Supplemental notes
Acid-Base Equilibria in Water
In the case below where water is the solvent:
X- + H3O+
HX + H2O
then:
H3O
X
Keq =
X
H2O
H
In this case concentration of H2 O = 55 M and is effectively unchanged since it is
present in such a large excess.
Then,
H3O
X
Ka = Keq [H2O] =
X
H
• Ka = dissociation constant and is a measure of acid strength.
• Larger Ka implies stronger acid.
• Range of Ka we may see is from 10-55 up to 107 ; 62 ! orders of magnitude.
Chemists use inverse log scale:
pKa = -logKa
the lower the pKa, the stronger the acid
pH = -log[H3 O+]
the lower the pH, more acidic the solution
Note for:
n
X
H
m
+
Keq
n-1
X
Y
Keq = 10-([pKa (HX)]-[pKa (HY)])
-10-
+
H
Y
m+1
Fall 2004 Supplemental notes
Examples
pKa values
HF
3
HF
H2O
16
NH3
33
+ OH
CH4
50-60
F
+ H2 O
pKa
3
16
Keq = 10-(3-16) = 10+13 So this reaction goes towards the right.
H2O
pKa
+ CH3
OH
16
+
CH4
~50
Keq = 10-(16-50) = 103 4 !!
So, pKa's are quantitative measures of acidity and allows one to make predictions
about reactions.
Example:
Conjgate Base:
CH3 - > NH2 - > OH- > :NH3 >
Acid
CH4 < :NH3 < H2 Ö: < NH4 + < H3 O+ <<<
pKa:
>50
33
16
10
H2 Ö: > >>>> HF
–2
H 2 F+
<< –10
Notice that both NH3 and H2 O can be both acid and base. Such compounds are
said to be amphoteric.
-11-
Fall 2004 Supplemental notes
Strengths of BrØnsted-Lowry Acids and Bases
Proton transfer reactions:
Proton transfer reactions can generate ions
hydroxide ion
H2O +
H2O
OH
+
H3O
hydronium ion
length of arrow indicates approximate position of equilibrium
Hydronium ion as acid:
H3O + NH3
H2O + NH4
Hydroxide ion as base:
OH
OH
OH
+
+
+
NH3
H2O +
NH2
CH4
H2O +
CH3
HF
H2O +
-12-
F
Fall 2004 Supplemental notes
Rules for Charge Stability of Ions with a Full Octet
Element effect
1.
Negative charge is most stable on most electronegative atom.
F- > R-O- > R2N- > R3CIncreasing Stability
2.
For atoms of similar electronegativity, the negative or positive charge is more
stable on the larger atom.
R-Te- > R-Se- > R-S- > R-OIncreasing Stability
R2SH+ > R2OH+
Increasing Stability
Why? Larger atoms distribute charge over a greater volume.
3.
Positive charge is most stable on least electronegative atom.
R3NH+ > R2OH+
Increasing Stability
These trends in stability can be used to predict directions of the acid-base reaction
shown above and others throughout the term.
LEARN THIS WELL!
-13-
Fall 2004 Supplemental notes
Periodic table of 246a
1
2
3
4
5
6
7
He
H
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Br
Kr
I
Xe
Electronegativity increases across row
Electron affinity increases across row
Acidity increases across row
Acidity of acids in a row
pKa
CH4 < NH3 < H2 O < HF
~50 ~32
16
3.5
Acidity of acids in a column
pKa
8
HF < HCl < HBr < HI
3.5
~–6
~–8 ~–10
-14-
Acidity increases down column
Bond strength to H decreases down column
Valence electrons
Fall 2004 Supplemental notes
F is more electronegative than I . So how do we explain this trend?


H-A → H+ + A:-
Consider reaction
We can use Hess law of summation to break up reaction into pieces
1. Bond breaking
H-A → H + A

2. Electron add to A

3. Ionization of H

Sum

e- + A → A:
H

→ H+ + e-
H-A → H+ + A-
-15-
Fall 2004 Supplemental notes
Now for comparison between acids:
1. Bond Breaking
a) Energy to break bonds drops dramatically down column
Bond
H-F
Bond Dissociation Energy (BDE) (kcal/mol): 136
H - Cl
103
H - Br
88
Reason: Lower orbital overlap
b) Energy to break bond doesn't vary so much across a row
Bond
CH3 -H
BDE (kcal/mol) 105
2. e- +A 
a)
NH2 -H
107
HO-H
119
F-H
136
→ A:- is electron affinity, EA
Electron affinity doesn't vary that much down a column
Atom
I
Br
Cl
F
EA(kcal/mol)
70
78
83
78
b)
Electron affinity increases dramatically across a row
Atom
EA(kcal/mol)
CH3 
1.8
NH3 
18
OH
42
F
78
3. Ionization of H to H+
This is the same for each acid so it doesn't enter into comparison.
-16-
H-I
71
Fall 2004 Supplemental notes
So,
For BDE
(kcal/mol)
HF
136
HI
71
∆BDE = 65 kcal/mol
F
78
I
70
∆EA = 8 kcal/mol
Favors HI
For EA
(kcal/mol)
Favors HF
But 65 >> 8 thus HI must be stronger acid than HF
So, down a column BDE dominates the strength of the acid.
In a row EA dominates the strength of the acid.
-17-
Fall 2004 Supplemental notes
Polar or Inductive Effects
Remember that opposite charges attract -- like charges repel
No relative stabilization
1)
2)
Charge spread over larger
volume-some stabilization
3)
Dipole- some significant
stablization
4)
Two dipoles- more significant
stablization
5)
Dipole oriented in wrong directiondestablization
6)
Remote dipole- weak stablization
All other things equal, if molecules have similar conjugate bases which experience
these environments then (ignoring entry #2!):
pKa: 5 > 1 > 6 > 3 > 4
-18-
Fall 2004 Supplemental notes
Examples
Number of polar groups:
OH
H
4.73
OH
H
Cl
OH
Cl
H
H
H
O
Cl
Cl
H
pKa
O
O
O
2.86
OH
Cl
Cl
1.26
0.064
Proximity of polar groups:
O
O
Cl
OH
OH
O
O
OH
OH
Cl
pKa
4.82
4.52
4.06
2.84
Cl
Resonance:
OH
O
OH
pKa
Why?
18
5
13 orders of magnitude
but also
O
O
O
O
O
O
less important
The minus charge is delocalized between the oxygens. The resonance structure on
the right inductively stabilizes oxygen (but is a minor contributor).
SUMMARY: Three major effects: Element effect (EA and BDE), Inductive effect, and
Resonance
-19-
Fall 2004 Supplemental notes
LEARN THESE VALUES
Conjugate
Acid
Conjugate
Base
pKa
HI
I-
-11
HBr
Br-
-8
HCl
Cl-
-6
H3O+
H2O
-2
SO3 -
SO3 H
O
F 3C
O
F 3C
-
OH
O
O
O
H 3C
-1
H 3C
OH
O
-
0
4.76
H2S
HS-
7.0
HCN
CN-
9.2
NH4+
NH3
9.2
R-SH
R-S-
10-12
R-OH
R-O-
16-20
NH3
NH2-
32
H2
H-
35
CH3↓
Increasing basicity:
DOWN in table
48
CH4
↑
Increasing acidity:
UP in table
A strong acid makes a weak base and vice versa.
-20-
Fall 2004 Supplemental notes
Hydrogen Bonding
A hydrogen bond is a particular type of a Lewis acid-Lewis base interaction.
• It can occur between a hydrogen atom attached to a heteroatom such as O,
F, N (called the hydrogen bond donor group) and an atom that has a lone
pair (typically also O, F, and N) the hydrogen bond acceptor.
• More generally any acidic hydrogen can be a hydrogen bond donor and any
Lewis base can be a hydrogen bond acceptor.
• Hydrogen bonding is a special case of dipole-dipole interactions, and it is
also an example of a weak covalent bonding interaction.
donor
acceptor
O
H
H
H
0.96Å
O
H
1.8-1.9Å
Note that the O–H---O angle is drawn to be 180°, I believe that this is the
preferred angle for hydrogen bonds.
-21-
Fall 2004 Supplemental notes
Effects of Hydrogen Bonding
• Hydrogen bonding affects the boiling point of solvents. Thus for water and
low molecular weight alcohols, the boiling points are unusually high since in
addition to overcoming van der Waals interactions, the hydrogen bonds
must be broken in order to vaporize the solvent.
• If such interactions did not occur it is likely that water would boil below
ambient temperature, which would make life on earth rather difficult.
• As we will see later, solvents capable of hydrogen bonding selectively
stabilize anions.
-22-
Fall 2004 Supplemental notes
Importance of Hydrogen Bonding to Life on Earth
Hydrogen bonds are critical to defining the base pairing in DNA. The
specificity of the hydrogen bonding interactions in DNA is thought to be
central to its ability to replicate with high fidelity.
H
O
-O
P
-O
N
N
H
O
O
O
H
HO
H
N
O
N
H
N
H
O-
P
O
N
O-
N
O
O
H H
H
H
H
OH
H H
Thymine
Adenine
H
-O
-O
N
O
P
O
O
H
N
O
O
H
HO
N
H
H
N
N
N
O
H
N
O
H
N
H
H H
O
H
H H
Guanine
Cytosine
-23-
H
H
OH
P
OO-
Fall 2004 Supplemental notes
Hydrogen Bonding and Proteins
Hydrogen bonds are critical to the so-called secondary structure of proteins (of
which enzymes are a subset).
• The primary structure is the sequence of amino acids that make up the
protein. The secondary structure is predominantly determined through
hydrogen bonding interactions. These interactions largely define the threedimensional structure of the protein.
• The actual sequence of amino acids determines what hydrogen bonds can be
formed. Much research is now devoted to understanding how to predict the
three-dimensional structure of proteins based upon the amino acid sequence.
• The three-dimensional structure of a protein determines its physical and
chemical properties.
• As an example, spider silk has a specific secondary structure (known as β
pleated sheets) that gives it strength in three dimensions (its strength per unit
weight is greater than that of steel!)
O
N
H
O
H
N
O
H
N
•
• The reactivity of an enzyme is defined by its three -dimensional structure.
-24-