Angewandte
Chemie
Table 1: Turnover frequencies (TOF) for the dehydration of xylose to
dehydration. Accordingly, this behavior is independent of the
furfural (reaction 1) in various solvents for homogeneous Brønsted acid
specific details of the acid-catalyzed mechanism, suggesting
that these solvent effects are related to solvation of the acidic-
proton catalyst.
[
a]
catalysts.
[
b]
ꢀ1
Solvent
Catalyst
T [K]
TOF [ks ]
1
We consider that these acid-catalyzed reactions are
controlled by two processes, the first of which is the
dissociation equilibrium of the acid, HB, into its acidic
H O
GVL
dioxane
SA
SA
SA
PSA
PSA
SA
SA
SA
448
448
448
448
448
418
418
418
1.5ꢂ0.05
46ꢂ2
2
21ꢂ1
+
ꢀ
H O
2.6ꢂ0.2
25ꢂ1
proton, H , and its conjugate base, B :
2
GVL
H O
0.090ꢂ0.01
5.1ꢂ0.2
1.2ꢂ0.09
HB Ð Hþ þ Bꢀ
2
ð1Þ
GVL
THF
In the second process, the acid-catalyzed conversion takes
place through a series of steps, one of which is assumed to be
the rate-determining step for simplicity. As shown elsewhere,
in this case, the overall rate can be written in terms of
a product of equilibrium constants for the steps that are not
rate-determining, multiplied by the equilibrium constant for
the formation of the transition state for the rate-determining
[a] Reaction conditions: xylose (0.15m), solvent (4 mL), and stirring at
7
00 rpm. [b] Organic solvents contained 10 wt% H O.
2
by 30–55 times compared to the values obtained when H O is
used as the solvent. Propylsulfonic acid (PSA), a weaker acid
than SA, exhibited a 10-fold increase in TOF1 for GVL
2
[9]
step. Accordingly, the overall rate can be expressed in terms
y
compared to the value in H O. Therefore, the reactivity of
of a single equilibrium constant (K
2
, which is a product of the
2
a homogeneous Brønsted acid catalyst is increased using
GVL as solvent, and it appears that this effect is more
pronounced for stronger homogeneous acids. More generally,
we have found that the high reactivity of a strong Brønsted
acid catalyst (i.e., SA) is also observed for other polar aprotic
solvents. For example, the results in Table 1 show that polar
aprotic solvents such as dioxane and THF also display
increased reactivity, similar to GVL.
aforementioned individual equilibrium constants) for the
formation of the rate-determining transition state from the
reactant R, as written below:
R þ Hþ Ð RHþy
ð2Þ
The rate of the second process is then given from
transition-state theory as:
The apparent activation energies for furfural formation
from xylose (reaction 1) and furfural degradation (reaction 2)
k T
h
B
y
þ
þ
r2 ¼
K ½Rꢁ½H ꢁ ¼ k ½Rꢁ½H ꢁ
ð3Þ
2
2
ꢀ
1
in H O were determined to be 145 and 85 kJmol , respec-
2
tively, which are in agreement with reported literature
where k is the Boltzmann constant, h is Planckꢀs constant,
B
[8]
values. The use of GVL as a solvent changes the energetics
of the reaction network for the conversion of xylose to favor
formation of the desired furfural product. The activation
energy barrier for xylose dehydration (reaction 1) is
and k is the overall rate constant for step 2.
2
+
The proton concentration [H ] can be written in terms of
the dissociation constant K1:
ꢀ1
pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi
decreased in GVL (114 kJmol ), whereas the barrier for
furfural degradation (reaction 2) is higher in GVL
2
ꢀ
K þ K þ 4K ½HB ꢁ
þ
1
1
1
o
ð4Þ
½
H ꢁ ¼
2
ꢀ1
(
105 kJmol ). Accordingly, furfural selectivities from
xylose of up to 75% can be achieved in GVL using SA as
the catalyst, compared to only 50% furfural selectivity from
where HB is the initial associated-acid concentration. The
rate of reaction per unit volume is then given by:
o
xylose in H O.
2
rate ¼ k
½Rꢁ½Hþꢁ
ð5Þ
Importantly, we have found that the increased reactivity
of a Brønsted acid catalyst in GVL is also observed for other
acid-catalyzed reactions, and this behavior is of general
significance. For example, the acid-catalyzed dehydration of
2
We suggest that the lower reactivity of a Brønsted acid
catalyst in water is caused in part by increased solvation of the
acidic proton by water molecules. For instance, the standard
Gibbs free energy change for solvation of a proton changes
from ꢀ1113 kJmol in liquid water to ꢀ1089 kJmol in an
aprotic solvent such as acetonitrile. Thus, the proton
catalyst is stabilized in water to a greater extent than in an
aprotic solvent (DG(H ) = 24 kJmol for this example),
and this proton stabilization would lead to lower reactivity in
water, provided that the solvent has a fractional effect
1
,2-propanediol to propanal displayed an 18-fold reactivity
increase using GVL as the solvent compared to the reaction in
ꢀ
1
ꢀ1
H O. Furthermore, the selectivity for propanal production
2
[
10]
increased from approximately 60% in H O to 75% in GVL.
2
Increases in reactivity were also achieved for a hydrolysis
reaction, which is the reverse of a dehydration reaction. The
acid-catalyzed hydrolysis of cellobiose to glucose showed
a 30-fold reactivity increase using GVL as the solvent
+
ꢀ1
solv
+
compared to the reaction in H O. These reactivity and
(f*DGsolv(H )) on the transition state for the acid-catalyzed
2
selectivity increases for 1,2-propanediol dehydration and
cellobiose hydrolysis are similar to those achieved for xylose
dehydration to furfural, and these results show that increased
reactivity using GVL as the solvent is not limited to xylose
reaction relative to stabilization of the reactant, as displayed
in Figure 1. Note that the solvent affects the stabilization of
+
ꢀ
both the proton (H ) and the conjugate base (B ), as well as
+
†
the protonated transition state (RH ).
Angew. Chem. Int. Ed. 2014, 53, 11872 –11875
ꢀ 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim