Lewis Acids and Bases (5.6 - 5.11)
A Lewis acid is an electron-pair acceptor
Gilbert N. Lewis
1875-1946
• A Lewis acid has an empty orbital (or can rearrange
electrons to create an empty orbital) that can “accept” an
electron pair to form a new bond
A Lewis base is an electron-pair donor
• A Lewis base has an accessible lone pair of electrons that
can be “donated” to a Lewis acid to form a new bond
The product of a reaction between a LA and LB is a complex
called an adduct, held together by a newly formed covalent
(dative) bond
From a MO perspective, we can consider the reactivity of frontier orbitals:
Empty (acceptor) orbital of a Lewis acid → LUMO
Filled (donor) orbital of a Lewis base → HOMO
Energy →
Adduct formation occurs via overlap:
• The stability of adducts comes from a
net lowering of energy of filled orbitals
LUMO
HOMO
⥯
⥯
LA
[LA←LB]
complex
or adduct
LB
• Creates new HOMO/LUMO pair
(bonding/antibonding MOs)
• The lower the energy of the empty
orbital on the LA (LUMO of LA), the
stronger the Lewis acid
• The higher the energy of the filled
orbital on the LB (HOMO of LB), the
stronger the Lewis base
F
B
F
The resulting overlap creates a new
covalent bond and a
Lewis Acid-Base Adduct
Lewis Acid
F
F
Lewis Base
H
N
H
H
The lone pair in the
N sp3 orbital can donate
into the empty 2p
orbital of the B
H
F
B
N
F
F
or
H
H
H
F
B
N
F
H
H
Bonds formed in a Lewis acid/base
reaction are referred to as
coordinate or dative bonds, and
are sometimes written as arrows
showing where the electrons came
from (the Lewis base)
Lewis Acid and Base Theory Explains Brønsted-Lowry
Lewis Acid
H+
The resulting overlap creates a new
covalent bond and a Lewis AcidBase Adduct
H
H
N
H
H
Lewis Base
The lone pair in the
N sp3 orbital can donate
into the empty 1s orbital
of a proton (H+)
H
N
H
H
•
Note connection to nucleophiles (e-pair
donor) and electrophiles (e-pair
acceptors) which are used to discuss
reaction rates (kinetics)
•
LA/LB highlight the equilibrium
(thermodynamic) properties of reactions
Examples of Lewis Acid-Base Chemistry
Identify the Lewis acid and Lewis base.
Determine the resulting LA-LB adduct.
(a) AlCl3 + NCl3 → ?
AlCl3, like BH3, has an incomplete octet and a vacant (3p)
orbital that can accept a pair of electrons ← LEWIS ACID
NCl3 has a full octet and a lone pair it can donate to form a
new covalent bond ← LEWIS BASE
Cl
Cl Al
Cl
Cl
+
Cl
N
Cl
Cl
Cl
Al
N
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Al
N
Cl
Cl
Examples of Lewis Acid-Base Chemistry
Identify the Lewis acid and Lewis base.
Determine the resulting LA-LB adduct.
(b) BCl3 + CH3OCH3 → ?
BCl3 has an incomplete octet and a vacant (2p) orbital that
can accept a pair of electrons ← LEWIS ACID
The O atom in dimethyl ether has a full octet and two lone
pairs it can donate to form a new covalent bond ← LEWIS BASE
Cl
Cl
B
Cl
Cl
+
H 3C
O
CH3
H 3C
B
Cl
Cl
Cl
O
CH3
H 3C
B
Cl
Cl
O
CH3
Examples of Lewis Acid-Base Chemistry
Identify the Lewis acid and Lewis base.
Determine the resulting LA-LB adduct.
(c) Fe3+ + 6H2O → ?
Fe3+ has vacant orbitals that can accept a pair of electrons
→ LEWIS ACID
The O atom in water has a full octet and two lone pairs it can
donate to form a new covalent bond ← LEWIS BASE
OH2
Fe3+
+ 6
H
O
H 2O
H
H 2O
Fe
OH2
3+
OH2
OH2
OH2
H 2O
H 2O
Fe
H 2O
3+
OH2
OH2
Equilibria in LA-LB Reactions and Adduct Formation
•
Reactions between free Lewis acids and Lewis bases involve the formation of
an adduct held together by a coordinate (= dative = covalent) bond (5.9a)
•
However, LA-LB reactions more typically involve equilibria with adducts
“trading partners”, and equilibrium constants depending on the
thermodynamics of the reaction
For example, displacement reactions (5.9b):
H 3C
Cl
O
H 3C
B
Cl
Cl
+
Cl
N
N
B
Cl
Cl
+
H 3C
O
CH3
And “double displacement” (a.k.a. metathesis) reactions (5.9c):
LA1-LB2 + LA2-LB2 ⇄ LA1-LB2 + LA2-LB2
“swapping partners…” (metathesis is Greek for “exchange”)
Examples of Lewis Acids (electron-pair acceptors) (5.6)
Need a vacant or “vacatable” energetically accessible orbital…
1. Molecules with incomplete octets of valence electrons
F
(a) trivalent Group 13 compounds: F B F + FF
F
B
→ good Lewis acids, so the adducts they
form tend to be relatively stable: AlX4-, BX4-…
→ used as Lewis-acid catalysts (e.g., Friedel-Crafts, 5.7b)
(b) carbocations or “carbenium” cations
e.g., triphenylmethyl, [C(C6H5)3]+
C
F
F
Examples of Lewis Acids (electron-pair acceptors) (5.6)
Need a vacant or “vacatable” energetically accesible orbital…
1. Molecules with incomplete octets of valence electrons
(c) a proton (H+) is a Lewis acid!
→ Lewis acid/base theory subsumes Brønsted-Lowry acid/base
theory!
+
H
Cl
→ a Brønsted-Lowry acid is Lewis acid-base adduct of the acid’s
conjugate base and a proton
→ dissociation of a Brønsted-Lowry acid in water is a
displacement reaction between two Lewis bases (conjugate
base of the acid and water) competing for a single Lewis acid (H+)
HA(aq) + H2O(l) ⇄ H3O+(aq) + A-(aq)
Examples of Lewis Acids (electron-pair acceptors) (5.6)
Need a vacant or “vacatable” energetically accessible orbital…
2. Metal cations can accept electron pairs from Lewis bases
(e.g., solvents) to form coordination complexes
→ aqua acids, solvated cations
OH2
Fe3+
+ 6
H
O
H 2O
H
H 2O
Fe
OH2
3+
OH2
OH2
OH2
H 2O
H 2O
Fe
H 2O
3+
OH2
OH2
Examples of Lewis Acids (electron-pair acceptors) (5.6)
Need a vacant or “vacatable” energetically accessible orbital…
3. Molecules with complete octets that can rearrange their
valence electrons to accept an electron pair
→ such molecules are said to be Lewis acidic
O C O + OH-
Similarly, SO2, SO3...
O
O
C
OH
Examples of Lewis Acids (electron-pair acceptors) (5.7)
Need a vacant or “vacatable” energetically accessible orbital…
4. Molecules that can “expand” their valence shell (or is
simply large enough!) to accept another electron pair
F
F
F
Si
F
F
+ 2 F-
F
F
Si
F
2-
F
F
→ tendency of Group 14 compounds (e.g., Group 14
tetrahalides) to act as Lewis acids by becoming 5/6 coordinate
e.g., SnCl4 + 2 Cl- → [SnCl6]2Note: lower valent Sn(II) halides can act as Lewis bases as
well as Lewis acids
Examples of Lewis Acids (electron-pair acceptors) (5.7)
Need a vacant or “vacatable” energetically accessible orbital…
4. Molecules that can “expand” their valence shell (or is
simply large enough!) to accept another electron pair
Group 14
C
Si
Ge
Sn
15
N
P
As
Sb
16
O
S
Se
Te
17
F
Cl
Br
I
→ Group 15 pentahalides (PnX5, Pn = N, P, As, Sb…) of the
heavier elements
F
Sb
X
X
X
Sb(III)
F
Sb
F + 2HF
F
F
Sb(V)
F
[H2F]+ F Sb F
F
F
F
-
Examples of Lewis Acids (electron-pair acceptors) (5.7)
Need a vacant or “vacatable” energetically accessible orbital…
5. Group 16 Lewis Acids
Sulfur dioxide (like CO2) can act as a Lewis acid (or a LB!)
O
O
S
O
O
S
O
+
N
R
R
R
R
S N R
R
O
Sulfur trioxide is a stronger Lewis acid (and much weaker LB)
O
O
S
O
O
O
S
O
O
+
N
R
R
R
R
O S N R
R
O
Examples of Lewis Acids (electron-pair acceptors) (5.7)
Need a vacant or “vacatable” energetically accessible orbital…
6. Group 17 diatomics
• I2 and Br2 are mildly Lewis acidic
• vacant σ* orbital is the “acceptor orbital”
• violet I2 turns brown in Lewis-basic solvents
I
I
+
I
I
I
I
triiodide anion (I3-)
I
I
I
Polyhalide anions (17.10)
• An I3- anion can similarly interact with other I2
molecules and form larger mononegative
polyiodides of composition [(I2)nI-]
• The triiodide anion is the most stable of the
series
• Cl3-, Br3-, BrI2- also known
Figure 17.12
Focusing on the 3 x pz orbitals from three I atoms, we can look at the formation of
the Lewis acid/base adduct:
• 3 AOs form 3 MOs
• Bonding, anti-bonding and non-bonding
• Bond order: 2 electrons in a bonding
MO shared by two bonds (1/2 bond
per pair of atoms)
I
Figure 17.11
I
I
I
I
I
I
I
I