Acids and Bases
Brønsted-Lowry Acids and Bases
• A Brønsted-Lowry acid is a proton donor.
• A Brønsted-Lowry base is a proton acceptor.
• H+ = proton
Figure 2.1
Examples of Brønsted–Lowry
acids and bases
1
Some molecules contain both hydrogen atoms and lone
pairs and thus, can act either as acids or bases,
depending on the particular reaction. An example is the
addictive pain reliever morphine.
2
Which of the following molecules is not a BromstedLowry acid?
HBr
NH3
CCl4
Bromsted-Lowry acid must
contain a proton.
Which of the following is not a Bromsted-Lowry base?
H3C
CH3
O
Bromsted-Lowry bases must have a lone pair or pi
bond to donate.
3
Reactions of Brønsted-Lowry Acids and Bases
• A Brønsted-Lowry acid base reaction results in the transfer of
a proton from an acid to a base.
• In an acid-base reaction, one bond is broken, and another
one is formed.
• The electron pair of the base B: forms a new bond to the
proton of the acid.
• The acid H—A loses a proton, leaving the electron pair in the
H—A bond on A.
4
• The movement of electrons in reactions can be illustrated using
curved arrow notation. Because two electron pairs are involved
in this reaction, two curved arrows are needed.
• Loss of a proton from an acid forms its conjugate base.
• Gain of a proton by a base forms its conjugate acid.
• A double reaction arrow is used between starting materials and
products to indicate that the reaction can proceed in the forward
and reverse directions. These are equilibrium arrows.
Examples:
5
Label the acid, base, conjugate acid and conjugate
base in the following reaction.
O
O
+
O
CH3
+
HO
H3C
H3C
CH3
CH2
CA
CB
B
A
CH3
Use curved arrows to show the movement of electron
pairs.
O
O
+
O
CH3
+
HO
H3C
CH2
H3C
CH3
CH2
H
6
Draw the products of the following reaction ( use curved
arrows to show movement of electron pairs)
H3C
H3C
H
H
+
NH2
NH2
HCl
H
+
H
H
Cl
+
+
H3C
NH3
+ Cl
H
H
H
+
H
H
7
Acid Strength and pKa
• Acid strength is the tendency of an acid to donate a proton.
• The more readily a compound donates a proton, the
stronger an acid it is.
• Acidity is measured by an equilibrium constant.
• When a Brønsted-Lowry acid H—A is dissolved in water, an
acid-base reaction occurs, and an equilibrium constant can
be written for the reaction.
8
Because the concentration of the solvent H2O is essentially constant,
the equation can be rearranged and a new equilibrium constant, called
the acidity constant, Ka, can be defined.
It is generally more convenient when describing acid strength to use
“pKa” values than Ka values.
9
10
Factors that Determine Acid Strength
•
Anything that stabilizes a conjugate base A:¯ makes the
starting acid H—A more acidic.
•
Four factors affect the acidity of H—A. These are:
Element effects
Inductive effects
Resonance effects
Hybridization effects
•
No matter which factor is discussed, the same procedure is
always followed. To compare the acidity of any two acids:
o Always draw the conjugate bases.
o Determine which conjugate base is more stable.
o The more stable the conjugate base, the more acidic the
acid.
11
Element Effects—Trends in the Periodic Table.
Across a row of the periodic table, the acidity of H—A
increases as the electronegativity of A increases.
Positive or negative charge is stabilized when it is spread
over a larger volume.
12
• Down a column of the periodic table, the acidity of H—A
increases as the size of A increases.
• Size, and not electronegativity, determines acidity down a
column.
• The acidity of H—A increases both left-to-right across a row
and down a column of the periodic table.
• Although four factors determine the overall acidity of a
particular hydrogen atom, element effects—the identity of
A—is the single most important factor in determining the
acidity of the H—A bond.
13
Factors that Determine Acid Strength—Inductive Effects
• An inductive effect is the pull of electron density
through  bonds caused by electronegativity differences between atoms.
• In the example below, when we compare the acidities
of ethanol and 2,2,2-trifluoroethanol, we note that the
latter is more acidic than the former.
14
• The reason for the increased acidity of 2,2,2trifluoroethanol is that the three electronegative
fluorine atoms stabilize the negatively charged
conjugate base.
15
• When electron density is pulled away from the negative
charge through  bonds by very electronegative atoms, it is
referred to as an electron withdrawing inductive effect.
• More electronegative atoms stabilize regions of high electron
density by an electron withdrawing inductive effect.
• The more electronegative the atom and the closer it is to the
site of the negative charge, the greater the effect.
• The acidity of H—A increases with the presence of electron
withdrawing groups in A.
16
Factors that Determine Acid Strength—Resonance Effects
• Resonance is a third factor that influences acidity.
• In the example below, when we compare the acidities of
ethanol and acetic acid, we note that the latter is more acidic
than the former.
• When the conjugate bases of the two species are compared, it
is evident that the conjugate base of acetic acid enjoys
resonance stabilization, whereas that of ethanol does not.
17
• Resonance delocalization makes CH3COO¯ more
stable than CH3CH2O¯, so CH3COOH is a stronger acid
than CH3CH2OH.
• The acidity of H—A increases when the conjugate
base A:¯ is resonance stabilized.
18
• Electrostatic potential plots of CH3CH2O¯ and
CH3COO¯ below indicate that the negative charge is
concentrated on a single O in CH3CH2O¯, but
delocalized over both of the O atoms in CH3COO¯.
19
Factors that Determine Acid Strength—Hybridization
Effects
• The final factor affecting the acidity of H—A is the
hybridization of A.
Let us consider the relative acidities of three different
compounds containing C—H bonds.
• The higher the percent of s-character of the hybrid
orbital, the closer the lone pair is held to the nucleus,
and the more stable the conjugate base.
20
21
Figure 2.5
Summary of the factors that
determine acidity
22
Calculate the pka’s and which is the stronger
acid?
OH
ka=
pka=
CH3
10-10
10-41
10
41
OH
23
Which is the strongest base?
H20 (pka= 15.7)
NH3 (pka= 38)
CH4 (pka= 58)
CH4
Draw the products and determine the direction of the
equilibrium.
CH4
pka= 50
+
OH
CH3
+
H2O
15.7
24
Which of these bases is strong enough to deprotonate
CH3COOH (pka= 4.8) ?
H
Cl
pka=
25
35
-7
Both acetylene and hydrogen will work
(must have higher pka)
Not using pka, which is the stronger acid?
H2S
HBr
HBr, b/c Br is further to the right and further down in
the periodic table.
25
Which is the stronger acid?
F2CHCH2COOH
OR
Cl2CH2COOH
F2CHCH2COOH is the most acidic b/c acids are
stabilized by electron withdrawing groups and F is more
electronegative than Cl
Which is the stronger acid and why?
H
H
H
B/c the C-H bond involves sp2
hybridized orbital instead of
sp3
26
Commonly Used Acids in Organic Chemistry
In addition to the familiar acids HCl, H2SO4 and HNO3, a
number of other acids are often used in organic
reactions. Two examples are acetic acid and p-toluenesulfonic acid (TsOH).
27
Commonly Used Bases in Organic Chemistry
Common strong bases used in organic reactions are
more varied in structure.
28
It should be noted that:
• Strong bases have weak conjugate acids with high pKa
values, usually > 12.
• Strong bases have a net negative charge, but not all
negatively charged species are strong bases. For example,
none of the halides F¯, Cl¯, Br¯, or I¯, is a strong base.
• Carbanions, negatively charged carbon atoms, are especially
strong bases. A common example is butyllithium.
• Two other weaker organic bases are triethylamine and
pyridine.
29
Lewis Acids and Bases
• The Lewis definition of acids and bases is more general
than the BrØnsted-Lowry definition.
• A Lewis acid is an electron pair acceptor.
• A Lewis base is an electron pair donor.
• Lewis bases are structurally the same as BrØnsted-Lowry
bases. Both have an available electron pair—a lone pair or
an electron pair in a  bond.
• A BrØnsted-Lowry base always donates this electron pair
to a proton, but a Lewis base donates this electron pair to
anything that is electron deficient.
30
• A Lewis acid must be able to accept an electron pair, but
there are many ways for this to occur.
• All BrØnsted-Lowry acids are also Lewis acids, but the
reverse is not necessarily true.
o Any species that is electron deficient and capable of
accepting an electron pair is also a Lewis acid.
• Common examples of Lewis acids (which are not BrØnstedLowry acids) include BF3 and AlCl3. These compounds
contain elements in group 3A of the periodic table that can
accept an electron pair because they do not have filled
valence shells of electrons.
31
• Any reaction in which one species donates an electron pair
to another species is a Lewis acid-base reaction.
• In a Lewis acid-base reaction, a Lewis base donates an
electron pair to a Lewis acid.
• Lewis acid-base reactions illustrate a general pattern in
organic chemistry. Electron-rich species react with electronpoor species.
• In the simplest Lewis acid-base reaction one bond is formed
and no bonds are broken. This is illustrated in the reaction of
BF3 with H2O. H2O donates an electron pair to BF3 to form a
new bond.
32
• A Lewis acid is also called an electrophile.
• When a Lewis base reacts with an electrophile other
than a proton, the Lewis base is also called a
nucleophile. In this example, BF3 is the electrophile and
H2O is the nucleophile.
electrophile
nucleophile
33
• Two other examples are shown below. Note that in each
reaction, the electron pair is not removed from the
Lewis base. Instead, it is donated to an atom of the
Lewis acid and one new covalent bond is formed.
34
• In some Lewis acid-base reactions, one bond is formed
and one bond is broken. To draw the products of these
reactions, keep in mind the following steps:
Always identify the Lewis acid and base first.
Draw a curved arrow from the electron pair of the
base to the electron-deficient atom of the acid.
Count electron pairs and break a bond when needed
to keep the correct number of valence electrons.
35
Consider the Lewis acid-base reaction between
cyclohexene and H—Cl. The BrØnsted-Lowry acid HCl
is also a Lewis acid, and cyclohexene, having a bond,
is the Lewis base.
36
• To draw the product of this reaction, the electron pair in
the  bond of the Lewis base forms a new bond to the
proton of the Lewis acid, generating a carbocation.
• The H—Cl bond must break, giving its two electrons to
Cl, forming Cl¯.
• Because two electron pairs are involved, two curved
arrows are needed.
37
Which is not a Lewis base?
NH3
Lewis base must have an electron pair to donate.
Which is not a Lewis acid?
Br
OH
CH
B
Br
Br
None, all are Lewis acids b/c they either contain an
unfilled valence shell or an acidic proton.
38
39