Prussian Blue/NaNO2 as an Efficient Reagent for the Nitration of Phenols in Aqueous Bisulfate and Acetonitrile Medium: Synthetic and Kinetic Study

Article abstract of DOI:10.1002/kin.21068

The reaction kinetics of Prussian blue (PB)/NaNO₂ initiated for the nitration of phenols by in aqueous bisulfate and acetonitrile medium indicated first-order dependence on [phenol], [NaNO₂], and [PB]. An increase in [KHSO₄] accelerated the rate of nitration under otherwise similar conditions. The rate of nitration was faster in the solvent of higher dielectric constant (D). Observed results were in accordance with Amis and Kirkwood plots [log k′ vs. (1/D) and [(D ? 1)/(2D + 1)]. These findings together with the linearity of plots, log k′ versus (vol% of acetonitrile (ACN)) and mole fraction of (nx) ACN, probably indicate the importance of both eloctrostatic and nonelctrostatic forces, solvent–solute interactions during nitration of phenols. Reaction rates accelerated with the introduction of electron-donating groups and retarded with electron-withdrawing groups, which are interpreted by Hammett's theory of linear free energy relationship. Hammett's reaction constant (ρ) is a fairly large negative (ρ < 0) value, indicating attack of an electrophile on the aromatic ring. Furthermore, an increase in temperature decreased the reaction constant (ρ) values. This trend was useful in obtaining isokinetic temperature (β) from Exner's plot of ρ versus 1/T. Observed β value (337.8 K) is above the experimental temperature range (303–323 K), indicating that the enthalpy factors are probably more important in controlling the reaction.

Full text of DOI:10.1002/kin.21068

Prussian Blue/NaNO2 as an
Efficient Reagent for the
Nitration of Phenols in
Aqueous Bisulfate and
Acetonitrile Medium:
Synthetic and Kinetic Study
PASNOORI SRINIVAS, MUPPIDI SURESH, K. C. RAJANNA, G. KRISHNAIAH
Department of Chemistry, Osmania University, Hyderabad, 500 007, India
Received 30 June 2016; revised 9 November 2016; accepted 12 December 2016
DOI 10.1002/kin.21068
Published online 23 January 2017 in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: The reaction kinetics of Prussian blue (PB)/NaNO2 initiated for the nitration of
phenols by in aqueous bisulfate and acetonitrile medium indicated first-order dependence on
[phenol], [NaNO2], and [PB]. An increase in [KHSO4] accelerated the rate of nitration under
otherwise similar conditions. The rate of nitration was faster in the solvent of higher dielectric
constant (D). Observed results were in accordance with Amis and Kirkwood plots [log kꢀ vs.
(1/D) and [(D − 1)/(2D + 1)]. These findings together with the linearity of plots, log kꢀ versus
(vol% of acetonitrile (ACN)) and mole fraction of (nx) ACN, probably indicate the importance
of both eloctrostatic and nonelctrostatic forces, solvent–solute interactions during nitration
of phenols. Reaction rates accelerated with the introduction of electron-donating groups and
retarded with electron-withdrawing groups, which are interpreted by Hammett’s theory of
linear free energy relationship. Hammett’s reaction constant (ρ) is a fairly large negative (ρ <
0) value, indicating attack of an electrophile on the aromatic ring. Furthermore, an increase
in temperature decreased the reaction constant (ρ) values. This trend was useful in obtaining
isokinetic temperature (β) from Exner’s plot of ρ versus 1/T. Observed β value (337.8 K) is
above the experimental temperature range (303–323 K), indicating that the enthalpy factors
ꢁC
are probably more important in controlling the reaction.
Chem Kinet 49: 209–218, 2017
2017 Wiley Periodicals, Inc. Int J
INTRODUCTION
Prussian blue (PB) pigment has got the distinction of
being the first molecular coordination compounds (fer-
ric ferrocyanide complex (Fe4[Fe(CN)6]3), which was
discovered by the German artist Diesbach in 1703 [1].
Correspondence to: K. C. Rajanna; e-mail: kcrajannaou@
yahoo.com.
Supporting Information is available in the online issue at
www.wileyonlinelibrary.com.
ꢁC
2017 Wiley Periodicals, Inc.

210
SRINIVAS ET AL.
For about three centuries, it is used as important ingre-
dient in the preparation [1,2] of paints, printing inks,
laundry dyes, etc. However, its structural, electronic,
and magnetic properties and other transition metal ana-
logues have been studied only in the past decade [2–7].
PB exists as water-soluble (KFe{Fe(CN)6}·xH2O) and
insoluble (Fe4{Fe(CN)6}·3yH2O) forms [8]. PB and
its analogues are used as efficient adsorbents in several
catalytic processes [9], as well as an efficient tool for
molecular separation [10] and for the removal of heavy
metal ions in wine production [11], which could be at-
tributed to their microporous character. PB exhibited
several electrochemical applications for battery build-
ing [12], electronic switching, and as electrochromic
devices [13]. PB has also been used as an efficient
catalyst for H2O2 decomposition [14], transesterifica-
tion of β-ketoesters [15], and epoxidation of styrene
in the presence of tert-butyl hydroperoxide [16]. Re-
cently, a team led by Jose´ Ramo´n Gala´n-Mascaro´s [17]
revealed that Co–Fe Prussian blue coordination poly-
mers could be effectively used for the oxidation of
water. Recently, Adhikamsetty and Jonnalagadda re-
ported the kinetics and mechanism of water-soluble
PB formation [18].
Nitration of aromatic compounds is one of the most
important electrophilic aromatic substitution reactions
to introduce a nitrogroup into an aromatic ring of
organic compound. It is used for the production of
large varieties of nitroaromatics, which accomplished
a broad spectrum of proven applications in the syn-
thesis of new/small-sized materials, dyes, pharmaceu-
ticals, perfumes, and plastics since the beginning of
the 20th century. Several nitroaromatics are also used
as solvents and chemical intermediates and precursors
in the manufacture of synthetic dyestuffs and other
chemicals. For a longer time, a mixture of concen-
trated sulfuric and nitric acids is used as a source for
the generation of nitronium ion species, which is the
active ingredient for nitration. However, a large amount
of acid mixture is discharged through laboratory and
industrial outlets, which is the root cause to trigger
environmental issues. Several groups of researchers all
over the world suggested different alternative protocols
to to overcome these issues [19–21].
observation on the formation of the nitronium ion from
freezing point studies. The dissolution of HNO3 in
100% H2SO4 at the freezing point produces four ions,
according to the following equilibrium:
HNO3 + H2SO4 ꢀ NO+2 + H3O+ + 2 HSO4−
The existence of nitronium ion (NO+2 ) species has
been further confirmed by x-ray crystallography and
Raman spectroscopic studies [24–26]. Research re-
ports of Kecki [29] and Miller et al. [30] furnished
excellent literature and a detailed account of the role of
acidity in the nitration of aromatic compounds from de-
tailed kinetic studies. In another review, Kulkarni [31]
highlighted four different approaches for continuous
flow of nitration with microreactors. Enthused by the
foregoing aspects on nitration reaction, we have also
taken up a detailed colorimetric kinetic study compris-
ing PB/NaNO2-initiated nitration of phenols in aque-
ous bisulfate and acetonitrile medium. Nevertheless,
we have recently reported PB as an ecofriendly cata-
lyst for selective nitration of organic compounds in the
presence of small amounts of HNO3 for the synthe-
sis of nitroaromatics under conventional and noncon-
ventional conditions [34], which is different from the
present kinetic investigation.
EXPERIMENTAL
General
Chemicals used in this study are of reagent grade,
which were purchased from Aldrich, Merck, Loba,
or Fluka (India). Laboratory-made distilled water was
further purified over permanganate and distilled twice,
whereas other organic solvents were fractionally dis-
tilled before use. Products of the reactions were char-
acterized by spectroscopic methods and physical data
such as melting/boiling points, infrared (IR),1H NMR,
and mass spectroscopic studies.
Kinetic Method of Following the Reaction
On the other hand, several kinetic and instrumen-
tation studies have been taken up to understand the
mechanism of nitration since the 1904 [22–33], which
provided ample of evidence for the formation of ni-
tronium ion (NO+2 ). The kinetic studies taken up by
Ingold’s research group [27] and Martinsen [28] con-
tributed enormously to understand the role of nitron-
ium (NO+2 ) and nitracidium (H2NO3+) ion species in
the mechanism of nitration under different conditions.
Gillespie’s research group [24] suggested a noteworthy
A reaction flask containing known amounts of aque-
ous PB and NaNO2 solutions and another flask con-
taining phenol (prepared in acetonitrile) along with
requisite amounts of KHSO4 were thermostated in a
constant temperature bath at a desired temperature.
The reaction was initiated by mixing both sets of solu-
tions thoroughly. In all the experiments, [NaNO2] was
taken large excess over [PB] ([NaNO2]0 ꢀ [PB]0) in
aqueous acetonitrile/KHSO4 medium, so that in situ
generated [NO+2 ] was equal to [PB]0. Black-coated
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

PRUSSIAN BLUE/NaNO2 AS AN EFFICIENT REAGENT FOR THE NITRATION OF PHENOLS
211
flasks from outside were used to prevent photochem-
ical effects. Aliquots of the reaction mixture were
withdrawn into a cuvette and placed in electronically
thermostated cell compartment of the spectrophotome-
ter. The PB content could be estimated from the
previously constructed calibration curve showing ab-
sorbance (A) versus [PB] at 430 nm. Absorbance values
were in agreement with each other with an accuracy
of 3% error. For the purpose of kinetic studies, ab-
sorbance of nitrate species produced during the course
of reaction at a given time is defined as (At), whereas
the absorbance at infinite time (at the end of the re-
action) is considered as Aꢁ, and A0, the absorbance
(if any) before the start of reaction. Thus (Aꢁ – At) is
proportional to (a – x), the concentration of reactant at
any given time, whereas (Aꢁ – A0) is proportional to
initial concentration of the reactant (a).
Figure 1 Plot of ln [(Aꢁ– A0)/(Aꢁ – At)] versus time (first-
order plot) at 318 K in 25% MeCN [phenol] = 0.05 mol/dm3;
[NaNO2] = 0.05 mol/dm3; [PB] = 0.001 mol/dm3;
[KHSO4] = 0.01 mol/dm3. [Color figure can be viewed at
wileyonlinelibrary.com]
RESULTS AND DISCUSSION
Kinetic Observations in the Present Study
General Procedure for Synthesis of
Nitrocompounds under Acid-Free Kinetic
Conditions
(i) The plots of ln [(Aꢁ – A0)/(Aꢁ – At)] versus time
were linear with positive slope passing through origin,
under the conditions: [NaNO2]0, [phenol]0 ꢂ [PB)]0
(Fig. 1). This observation reveals first-order kinetics in
[PB]. First-order rate constant (kꢀ) could be obtained
from the slopes of these linear plots. Typical first-order
rate constant data are presented in Table I. (ii) However,
under pseudo–second-order conditions ([NaNO2]0 ꢂ
[PB)]0 = [S]0), plots of [1/(a – x)] or (1/(Aꢁ – At))
versus time have been found to be linear with a pos-
itive gradient and a definite intercept on the ordinate
(vertical axis) indicating overall second-order kinetics
(Fig. 2).
From the forgoing kinetic observations, it is clear
that order in [S] is also one, because order with re-
spect to [PB] is verified as one under pseudocondi-
tions (Figs. 1 and 2), and overall order is second order.
Second-order rate constant values for nitration of phe-
nols are compiled in Table II.
For the purpose of product analysis, a separate reaction
has been done under kinetic conditions, and progress
of the reaction was monitored by TLC. After comple-
tion, the reaction mixture was treated with 5% sodium
thiosulfate (hypo) solution. The organic layer was sep-
arated and dried over Na2SO4 and evaporated under
vacuum. Column chromatography was used to purify
crude product. A binary mixture of ethyl acetate and
hexane (3:7) was used as an eluent to get a pure product
4-nitrophenol as a yellow powder (mp 111–113°C) in
85% yield, IR (KBr) (νmax/cm−1): 3331, 1614, 1592,
1500, 1346; 1H NMR (300 MHz, CDCl3): δ 9.95
(s 1H, OH), 6.95 (d 2H, J = 8Hz), 8.15 (d 2H, J =
8Hz); m/z = 139.
PB/NaNO2/KHSO4-initiated reactions with several
other phenols also afforded the corresponding nitro-
phenol derivatives (Scheme 1) in about 1–3 h under
conventional stirred conditions at room temperature.
All the products were characterized by physical data,
1H NMR, and mass spectra, with authentic samples
are found to be satisfactory and in consonance with
our earlier reports [34].
Computation of Activation Parameters
Rate constants increased substantially within an in-
crease in the temperature. According to Eyring’s theory
OH
OH
Prussian Blue/NaNO2 / KHSO4 /(MeCN)
Stirring at Room temperature
NO2
R
R
Where R = EWD or ED group
Scheme 1 Nitration of phenols under different conditions.
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

212
SRINIVAS ET AL.
Table I Pseudo–First-Order Rate Constants (kꢀ) for Nitration of Phenols in 25% MeCN Medium
[Substrate] = 0.0125 mol/dm3; [NaNO2] = 0.125 mol/dm3; [KHSO4] = 0.0025 mol/dm3
104 [PB] = 2.50 mol/dm3; Acetonitrile (%) = 25 (V/V) ; Temperature = 300 K
Substrate
Phenol
Temperature (K)
Equation and R2
kꢀ (min)
300
313
323
300
313
323
300
313
323
300
313
323
300
313
323
313
318
323
300
313
323
300
313
323
300
313
323
y = 0.010x – 0.028; R2 = 0.991
y = 0.015x + 0.037; R2 = 0.990
y = 0.028x – 0.088; R2 = 0.989
y = 0.015x – 0.036; R2 = 0.989
y = 0.018x – 0.054; R2 = 0.987
y = 0.032x – 0.094; R2 = 0.977
y = 0.007x – 0.006; R2 = 0.990
y = 0.011x – 0.011; R2 = 0.990
y = 0.024x – 0.024; R2 = 0.993
y = 0.014x + 0.031; R2 = 0.992
y = 0.016x + 0.0217; R2 = 0.916
y = 0.029x + 0.046; R2 = 0.976
y = 0.007x; R2 = 0.996
0.010
0.015
0.028
0.015
0.018
0.032
0.007
0.011
0.024
0.014
0.016
0.029
0.007
0.019
0.026
0.012
0.015
0.019
0.008
0.013
0.025
0.005
0.007
0.101
0.005
0.009
0.021
p-Cresol
p-Cl phenol
m-Cresol
Beta napthol
p-nitrophenol
Resorcinol
Catechol
y = 0.019x – 0.030; R2 = 0.976
y = 0.026x + 0.0230; R2 = 0.973
y = 0.012x + 0.027; R2 = 0.927
y = 0.015x + 0.123; R2 = 0.966
y = 0.019x + 0.025; R2 = 0.968
y = 0.008x – 0.018; R2 = 0.995
y = 0.013x – 0.034; R2 = 0.994
y = 0.025x – 0.141; R2 = 0.953
y = 0.005x + 0.024; R2 = 0.986
y = 0.007x + 0.023; R2 = 0.991
y = 0.101x – 0.0182; R2 = 0.993
y = 0.005x + 0.024; R2 = 0.986
y = 0.009x + 0.024; R2 = 0.993
y = 0.021x – 0.008; R2 = 0.998
m-Chlorophenol
of reaction rates [35], free energy of activation (࢞G#)
can be correlated to the rate constant (k) at any given
temperature,
where the transmission coefficient (kt) is equal to
unity; R, N, h, and T represent the gas constant,
Avogadro number, Planck’s constant, and reaction tem-
perature, respectively; and ࢞G# represents free energy
of activation. Rearrangement of the above equation
gives
ꢀ
ꢁ
k = (kt ) (RT /Nh) exp −ꢀG#/RT
(1)
ꢀ
ꢁ
(Nhk/RT ) = exp −ꢀG#/RT
(2)
Natural logarithms for the above equation gives
ꢀ
ꢁ
ln (Nhk/RT ) = − ꢀG#/RT
(3)
ꢀ
ꢁ ꢀ
ꢀ
ꢁꢁ
Rearranging for ꢀG# : ꢀG# = −RT ln Nhk/RT
(4)
Equation (4) is used to calculate the free energy
of activation (࢞G#) at various temperatures, and the
Gibbs–Helmholtz equation (5) for the evaluation of
enthalpy and entropies of activation (࢞H# and ࢞S#)
and the representative plot (Fig. 3):
Figure 2 Plot of [1/(Aꢁ – At)] versus time (second-order
plot) in 25% MeCN [phenol] = 0.001 mol/dm3; [NaNO2] =
0.1 mol/dm3; [PB]
=
0.001 mol/dm3; [KHSO4]
=
0.1 mol/dm3. [Color figure can be viewed at wileyonlineli-
brary.com]
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

PRUSSIAN BLUE/NaNO2 AS AN EFFICIENT REAGENT FOR THE NITRATION OF PHENOLS
213
Table II Second-Order Rate Constants and Activation Parameters for Nitration of Phenols
ࣔ
ࣔ
ࣔ
Temperature
(K)
࢞G
࢞H
−࢞S
Substrate
Phenol
k
(kJ/mol)
Equation and R2
(kJ/mol)
(J/K/mol)
135
300
313
318
323
300
313
318
323
300
313
318
323
300
313
318
323
300
313
318
323
313
318
323
300
313
323
300
313
318
323
300
313
318
323
0.8
1.2
1.76
2.24
1.2
74.05536
76.36342
76.56817
77.16634
73.04405
75.84523
76.02672
76.80776
74.94498
77.12679
77.09871
77.5803
73.21613
76.15173
76.33812
77.07211
74.94498
75.70453
76.45064
77.36535
76.90036
77.58074
78.20766
74.61192
76.69207
77.47068
74.31815
76.69207
77.41011
77.93889
75.7842
y = 0.135x + 33.55; R2 = 0.962
33.55
p-Cresol
y = 0.163x + 24.31; R2 = 0.959
y = 0.114x + 40.69; R2 = 0.920
y = 0.168x + 23.03; R2 = 0.954
y = 0.100x + 44.67; R2 = 0.917
24.31
40.69
23.03
44.67
163
114
168
100
1.44
2.16
2.56
0.56
0.88
1.44
1.92
1.12
1.28
1.92
2.32
0.56
1.52
1.84
2.08
0.96
1.2
1.52
0.64
1.04
2
0.72
1.04
1.28
1.68
0.4
p-Cl Phenol
m-Cresol
Beta napthol
p-nitrophenol
Resorcinol
Catechol
y = 0.056x – 16.58; R2 = 0.993
y = 0.123x + 37.82; R2 = 0.963
y = 0.160x + 26.16; R2 = 0.988
16.58
37.82
26.16
56
123
160
m-Chlorophenol
y = 0.094x + 47.68; R2 = 0.900
47.68
94
0.72
1.2
1.68
77.64899
77.58074
77.93889
ꢀG# = ꢀH # − TꢀS#
(5)
Effect of variation of [KHSO4]. In the absence of
KHSO4, the reaction did not proceed even under reflux
temperatures. Reactions triggered only in the presence
of KHSO4, and rate of the reaction enhanced with an
increase in KHSO4 concentration. The plots of kꢀ ver-
sus [KHSO4] were found linear passing through origin
(Fig. 4), establishing the role of HSO4− in the genera-
tion of the active nitronium ion.
Figure 3 Gibbs–Helmholtz plot of ࢞G# versus T for nitra-
tion of p-nitrophenol. [Color figure can be viewed at wiley-
onlinelibrary.com]
Reactive Species and Mechanism
Nature of the solvent and its properties are known
to play an important role to influencing the reaction
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

214
SRINIVAS ET AL.
Figure 4 Plot of kꢀ versus [KHSO4]. [Color figure can be
viewed at wileyonlinelibrary.com]
Figure 5 Plot of log k versus 1/D. [Color figure can be
viewed at wileyonlinelibrary.com]
rates with or without influencing the reaction mecha-
nism [36–42]. Kinetics of the nitration has been stud-
ied in different dielectric media using binary solvent
mixtures of acetonitrile and water of different com-
positions, and the rate data are presented in Table III.
This affords a change in the dielectric constant (D) of
the medium. The authors tried to interpret the observed
kinetic results using semiquantitative relationships de-
veloped by Amis, Laidler-Eyring, Kirkwood, Hughes,
Ingold, and others [39,40], which are mainly concerned
with the rate of the reaction (or rate constant) as a
function of bulk dielectric constant (D). Accordingly,
the plots of log kꢀ versus 1/D has been found linear
with negative slope (Fig. 5), whereas the plot of log
kꢀ versus [(D − 1)/(2D + 1)] was linear with positive
slope (Fig. 6). These observations probably indicate
the participation of anionic species and molecular or
(dipolar) species in the rate-determining step. These
observations may probably indicate the participation
of anionic and molecular species in the rate-limiting
step.
were also linear with negative slopes (Figs. 7 and 8).
These observations probably indicate that apart from
the dielectric constant, solvent–solute interactions, and
other solvent properties might be important in control-
ling the reaction rates. This aspect could be supported
from the suggestions of Robinson and Stokes [37]
and findings of Brown and Hudson [38]. According
to Robinson and Stokes [37], the bulk dielectric con-
stant is entirely different from internal dielectric con-
stant especially in the solvent of low dielectric media.
Brown and Hudson have found that the rate of hydrol-
ysis of acid chlorides in DMF is almost equal to 95%
dioxane in spite of the difference in their dielectric
constants [38]. A perusal of literature depicts several
such anomalies [41].
UV–Visible Spectroscopic Studies. Since the reac-
tions are conducted in aqueous acetonitrile medium,
sodium nitrite (NaNO2) and KHSO4 being strong ele−c-
On the other hand, the plots of log kꢀꢀ versus volume
percentage (vol%) and mole fraction (nx) of acetonitrile
trolytes undergo dissociation to afford nitrite (NO2
)
Table III Effect of Temperature/Solvent on Hammett’s Reaction Constant (ρ)
Hammett’s equation: log k = log k0 ρσ
Rate Constant (k) at Given Temperature (T) (K)
Substrate
Hammett’s σ
300
313
318
323
Phenol
p-Cresol
p-Chloro
m-Cresol
Resorcinol
m-Chlorophenol
0
0.8
1.2
0.56
1.12
0.64
0.4
1.2
1.44
0.88
1.28
1.04
1.76
2.16
1.44
1.92
1.60
2.24
2.56
1.92
2.32
2.00
1.68
−0.17
0.23
−0.07
0.12
0.37
0.72
1.20
Hammett’s ρ (slope of Hammett’s plot)
Equation for the Exner’s plot: ρ versus 1/T
Isokinetic temperature (β) value
−0.903
−0.556
−0.455
−0.325
ρ = A [1 − β/T] = A− (Aβ)/T y = 2417x − 7.155; R2 = 0.998
337.8 K
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

PRUSSIAN BLUE/NaNO2 AS AN EFFICIENT REAGENT FOR THE NITRATION OF PHENOLS
215
Figure 9 PB in water.
Figure 6 Plot of log k versus (D – 1)/(2D + 1). [Color
figure can be viewed at wileyonlinelibrary.com]
Figure 10 PB in aqueous ACN.
and bisulfate (HSO4−). To have an insight into the in-
teraction between PB with anionic species such as ni-
trite (NO2−), bisulfate (HSO4−), and molecular species
such as acetonitrile, we have taken up an elaborate
UV–visible spectroscopic study of PB under differ-
ent conditions and results are presented in Figs. 9–12,
respectively.
Figure 7 Plot of log k versus MeCN%(v/v). [Color figure
can be viewed at wileyonlinelibrary.com]
UV–visible spectrum of aqueous PB (Fig. 9) indi-
cated two bands, one hump at 430 and the other broad
band at 675–680 nm. This observation is in close agree-
ment with that described by Jonnalagadda and Ad-
hikamsetty [18]. When the spectrum of PB solution is
scanned in aqueous acetonitrile, both the above-cited
peaks underwent hypochromic shift (Fig. 10), indicat-
ingsolvent (water)–solvent (ACN) andsolvent (ACN)–
solute(PB) interactions. Theseinteractions might bring
about some changes in the coordination spheres of Fe
present in PB. When the spectrum of PB is scanned
in aqueous acetonitrile medium containing KHSO4,
peak observed at 430 underwent a blueshift by 15 nm,
while another broad band underwent a redshift from at
Figure 8 Plot of log k versus mole fraction (nx). [Color
figure can be viewed at wileyonlinelibrary.com]
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

216
SRINIVAS ET AL.
nitration was enhanced with an increase in [KHSO4].
Observed rate acceleration could be attributed to the
available proton due to the dissociation of bisulfate
into proton and sulfate ions according to the following
equilibrium:
HSO−4 ꢀ H+ + SO24−
(6)
Soluble-PB anion (Fe{Fe((CN)6}−) then trans-
forms into protonated-PB species ([HFe{Fe((CN)6}]),
as shown in the following equilibrium:
ꢂ
ꢃ
Fe {Fe((CN)6}− + H+ ꢀ HFe {Fe((CN)6}
(7)
(Soluble PB anion)
(Protonated PB)
−
Figure 11 PB-KHSO4 in aqueous ACN.
On the other hand, an interaction between and
Fe(III)-enriched PB (hard acid) and SO 24− (hard base)
is also more likely according to Pearson’s Hard and
soft acids and bases (HSAB) theory, because hard acid
prefers to combine with hard base [42]. These forego-
ing kinetic features coupled with UV–visible spectro-
scopic results substantiate a plausible stepwise mech-
anism with the participation of [KHFeSO4 (NO2) [Fe
((CN)6)2−]] and molecular species (RC6H4X) in the
slow step as shown in Scheme 2.
675–680 nm to 680–700 nm (Fig. 11). This observa-
tion could be attributed to a change in the coordination
spheres of PB due to PB with the interaction of bisul-
fate ions. Finally, when the spectrum of PB is scanned
in aqueous acetonitrile medium containing KHSO4 and
NaNO2 peak observed at 430 underwent a blueshift by
10 nm, whereas another broad band underwent a red-
shift from 675–680 to 675–700 nm followed by slight
hyperchromic shift (Fig. 12). These changes could
most probably suggest the interaction of PB with nitrite
as well as with bisulfate. Furthermore, it is of interest to
note that protonation of ferrocyanide ion as such is too
fast and the H[Fe((CN)6]3− ion is known to react with
other species [18], but with its first protonation constant
pK = 4.19, i.e., K = 6.45 × 10−5 [18]. Thus an increase
in acid concentration could also lead to higher concen-
tration of protonated PB [HFe{Fe((CN)6}] due to the
reaction between H+ and Fe{Fe((CN)6}− ion (anionic
form of soluble PB). In the present study, the rate of
Quantitative Structure Reactivity Study
An insight into the kinetic data of nitration of phenols
revealed that the rates accelerated with the introduction
of electron-donating groups and retarded with electron-
withdrawing groups and found to follow the following
sequence: p-Me > m-Me > H > o-OH > m-OH >
m-Cl.
Accordingly, Hammett’s linear free energy rela-
tionships plots were tried to interpret these results.
Hammett’s log (k/k0) versus σ plot (Fig. 13) exhibited
very good linear relationship as evidenced from the
correlation coefficient values (R2), which were greater
Figure 13 Typical Hammett’s plot at 318 K. [Color figure
can be viewed at wileyonlinelibrary.com]
−
Figure 12 PB-KHSO4 NaNO2 in aqueous ACN.
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

PRUSSIAN BLUE/NaNO2 AS AN EFFICIENT REAGENT FOR THE NITRATION OF PHENOLS
217
KHFe(SO4)[Fe(CN)6]-
-
+ HSO4
KFe[Fe(CN)6]
PB (soluble)
PB (mono anionic form)
KHFe(SO4)[Fe(CN)6]-
PB (mono anionic form)
KHFe(SO4)(NO2)[Fe(CN)6]2-
PB (di anionic form)
-
+ NO2
+
X
NO2
ACN(Solvent)
KHFe(SO4)(NO2)[Fe(CN)6]2-
PB (di anionic form)
+ RC6H4X
Slow
R
KFe(ACN)2[Fe(CN)6]2- + HSO4
Solvated PB (di anionic form)
-
X
X
H+
R
NO2
NO2
R
+
Scheme 2 Plausible mechanism for nitration of phenols.
than 0.960 [43]:
importance of greater solvation in the transition state.
Put together, it appears that both entropy and enthalpy
values are important in this study.
ꢀ
ꢁ
log k/k0 = σρ
(8)
The negative value of ρ indicates that the electron
flow away from the aromatic ring takes place in the
rate-determining step by producing electron-deficient
transition state (often a positive charged activated com-
plex). The rate-determining step for most electrophilic
substitutions shows the electrophile reacts with the aro-
matic ring to form the arenium ion. Data presented
in Table III revealed that the rho (ρ) values obtained
from the present experiments are fairly large negative
values (ρ < 0), indicating attack of an electrophile on
the aromatic ring.
CONCLUSIONS
In summary, we have developed an acid-free proto-
col for nitration of phenols with PB/NaNO2 in an
aqueous bisulfate and acetonitrile medium. The re-
action followed second-order kinetics with first-order
dependence on [PB] and [phenol] when [NaNO2] is
taken large excess over [PB]. Rates of nitration (log k)
were found to fit into either Amis or Kirkwood plots
[log kꢀꢀ vs. (1/D) or [(D – 1)/(2D + 1)]. In addition to
this, the observed results may also indicate the impor-
tance of equilibrium as well as frictional solvent effects
and solvent–solute interactions for solvation of transi-
tion state during nitration of phenols. Reaction rates
accelerated with the introduction of electron-donating
groups and retarded with electron-withdrawing groups
with the following sequence: p-Me > m-Me > H >
o-OH > m-OH > m-Cl. The results are interpreted by
Hammett’s theory of linear free energy relationship.
The reaction constant (Hammett’s ρ) values obtained
from the present experiments are fairly large negative
values (ρ < 0), indicating attack of an electrophile
on the aromatic ring. An increase in temperature de-
creases the reaction constant (ρ) values, and Exner’s
plot of ρ versus 1/T gave isokinetic temperature (β)
with a value of 337 K, which is slightly higher than
the experimental temperature range (303–323 K). This
observation could probably suggest that enthalpy fac-
tors are important in controlling the reaction. Large
According to Exner [44,45], ρ values for a given
reaction are influenced by the temperature according
to the following relation:
(ρ) = A [1−β/T ] = A− Aβ/T
where A is a constant and β is the isokinetic tempera-
ture. When β = T, ρ = 0, thus isokinetic temperature
is the temperature at which the effect of substituent
on the rate of reaction vanishes and all the substituted
compounds in a given series have the same reactiv-
ity. In the present study, the plot of rho (ρ) versus
(1/T) gave a linear plot with slope of (A β), and in-
tercept was equal to A. The obtained (slope/intercept)
ratio afforded the isokinetic temperature (β) with a
value of 338 K, which is above the experimental
temperature range (303–323 K), suggesting that en-
thalpy factors are important in controlling the reaction.
Large negative entropy values probably indicate the
International Journal of Chemical Kinetics DOI 10.1002/kin.21068

218
SRINIVAS ET AL.
negative entropy values probably indicate the impor-
tance of greater solvation in the transition state.
22. (a) Ri, T.; Eyring, H. J Chem Phys 1940, 8, 433;
(b) Walden, P. Angew Chem 1924, 37, 390.
23. (a) Cox, E. G.; Jeffery, G. A.; Truter, M. R. Nature 1948,
162, 258; (b) Grison, E.; Ericks, K.; de Vries, J. L. Acta
Cryst 1950, 3, 290; (c) Goddard, D. R.; Hughes, E. D.;
Ingold, C. K. J Chem Soc 1950, 1, 2559.
BIBLIOGRAPHY
1. Woodward, J. Philos Trans 1724, 33, 15–17.
2. (a) Dunbar, K. R.; Heintz, R. A. Prog Inorg Chem
1997, 45, 283; (b) Shatruk, M.; Dragulescu-Andrasi,
A.; Chambers, K. E.; Stoian, S. A.; Bominaar, E. L.;
Catalina Achim, C.; Kim, R.; Dunbar, K. E. J Am Chem
Soc 2007, 129, 6104–6116.
3. Robin, M. B. Inorg Chem 1962, 1, 337–342.
4. Sano, H.; Hashimoto, F. Bull Chem Soc Jpn 1965, 38,
684–685.
24. (a) Gillespie, R. J.; Graham, J.; Hughes, E. D.; Ingold,
C. K.; Peeling, E. R. Nature 1946, 158, 480; (b) Gille-
spie, R. J.; Millen, D. J. Q Rev (London) 1948, 2, 277.
25. Chedin, J. Ann Chim 1937, 8, 295.
26. Olah, G. A.; Kuhn, S. J.; Flood, S. H. J Am Chem Soc
1961, 83, 4581–4585 and references cited there in.
27. (a) Ingold, C. K.; Hughes, E. D.; Reed, R. I. Nature
1946, 158, 448; (b) Halberstadt, E. S.; Hughes, E. D.;
Ingold, C. K. Nature 1946, 158, 514; (c) Halberstadt,
E. S.; Hughes, E. D.; Ingold, C. K. J Chem Soc 1950,
2441.
5. Walker, R. G.; Watkins, K. O. Inorg Chem 1968, 7,
885–888.
6. Ito, A.; Suenaga, M.; Ono, K. J Chem Phys 1968, 48,
3597–3599.
28. Martinsen, H. Z Phys Chem 1904, 50, 385; 1907, 59,
605.
7. Mayoh, B.; Day, P. J Chem Soc, Dalton Trans 1976,
1483–1486.
8. Balmaseda, J.; Reguera, E.; Fernandez, J.; Gordillo, A.;
Yee-Madeira, H. J. Phys Chem Solids 2003, 64, 685.
9. Reguera, E.; Fernandez, J.; Duque J. Polyhedron 1994,
13, 479.
10. Balmaseda, J.; Reguera, E.; Gomez, A.; Diaz, B.; Autie,
M. Microporous Mesoporous Mater 2002, 54, 285.
11. Wenker, D.; Spiess, B.; Laugel, P.; Lapp, C. Food Addit
Contam 1989, 6, 351.
12. Neff, V. D. J Electrochem Soc 1985, 132, 1382.
13. (a) DeBerry, D. W.; Viehbeck, A. J Electrochem Soc
1983, 130, 249; (b) Gadet, V.; Mallah, T.; Castro, I.;
Verdaguer, M. J Am Chem Soc 1992, 114, 9213.
14. Uyanik, G.; Pekin, B. J Catal 1970, 19, 195–203.
15. Srinivas, P.; Rajanna, K. C.; Krishnaiah, G.; Satish
Kumar, M.; Narender Reddy, J. Synth React Inorg Met-
Org Nano-Met Chem 2014, 44, 1212–1220.
16. Liang, Y.; Yi, C.; Tricard, S.; Fang, J.; Zhao, J.; Shen,
W. RSC Adv 2015, 5, 17993–17999.
17. Pintado, S.; Goberna-Ferro´n, S.; Escudero-Ada´n, E. C.;
Gala´n-Mascaro´s, J. R. J Am Chem Soc 2013, 135,
13270–13273.
18. Adhikamsetty, R. K.; Jonnalagadda, S. B. Bull Chem
Soc Ethiop 2009, 23, 47–54.
19. (a) Booth, G. In Ullmann’s Encyclopedia of Industrial
Chemistry; Wiley-VCH: Weinheim, Germany 2005;
(b) Olah, G. A.; Malhotra, R.; Narang, S. C. Nitration
Methods and Mechanisms (Organic Nitro Chemistry Se-
ries), Vol. A17; VCH: New York, 1989, p. 411–455
(1991).
20. (a) Hoggett, J. G.; Monodie, R. B.; Penton, J. R.;
Schofield, K. Nitration and Aromatic Reactivity; Cam-
bridge University Press: London (UK), 1971; (b) Ono,
N. The Nitro Group in Organic Synthesis; Wiley-VCH:
New York, 2001.
29. Kecki, Z. Tetrahedron 1961, 12, 23.
30. Miller, R. C.; Noyce, D. S.; Vermeulen, T. Ind Eng Chem
1964, 56, 43.
31. Amol, A.; Kulkarni, B. J Org Chem 2014, 10, 405.
32. (a) Ramesh, K.; Shylaja, S.; Rajanna, K. C.; Giridhar
Reddy, P.; Saiprakash, P. K. Adv Phys Chem 2013,
146585; (b) Int J Chem Kinet 2014, 46, 126.
33. Satish Kumar, M.; Rajanna, K. C.; Venkateswarlu, M.;
Lakshman Rao, K. Int J Chem Kinet 2016, 48, 171–196.
34. Srinivas, P.; Rajanna, K. C.; Satish Kumar, M.;
Rajendar Reddy, K. Synth React Inorg Metal-Org Nano-
Met Chem 2014, 44, 364.
35. (a) Glasstone, S.; Laidier, K. J.; Eyring, H. Theory
of Rate Processes; McGraw-Hil: New York, 1961;
(b) Laidler, K. J. Chemical Kinetics; Pearson Educa-
tion: Singapore, 2004.
36. Christian, R. Solvents and Solvent Effects in Organic
Chemistry; Wiley-VCH: Weinheim, Germany, 2003.
37. Robinson, R. A.; Stokes, R. H. Electrolyte Solutions;
Butterworths Scientific Publications: London, 1955.
38. Brown, D. A.; Hudson, R. F. J Chem Soc 1953, 883.
39. (a) Amis, E. S. Solvent Effects on Reaction Rates
and Mechanisms; Academic Press: London, 1964;
(b) Kirkwood, J. G. J Chem Phys 1924, 2, 251.
40. (a) Ingold, C. K. Structure and Mechanism in Organic
Chemistry; Bell: London, 1953; (b) Hughes, E. D.; In-
gold, C. K. J Chem Soc 1935, 244.
41. Shorter, J. Correlation Analysis in Organic Chem-
istry: Introduction to Linear Free-energy Relationships;
Clarendon Press: Oxford, UK, 1973, pp. 70.
42. Basolo, F.; Pearson, R. G. Mechanisms of Inorganic
Reactions—A Study of Metal Complexes in Solution,
2nd ed.; Wiley : New York, 1967.
43. (a) Hammett, L. P. Physical Organic Chemistry;
Mc-Graw-Hill: New York, 1940; (b) Wiberg, K. B. Phys-
ical Organic Chemistry; Wiley: New York, 1964.
44. Exner, O. Coll Czech Chem Commun 1955, 29, 1094.
45. (a) Petersen, J. R. C. J Org Chem 1964, 29, 3133;
(b) Liu, L.; Guo, Q. X. Chem Rev 2001, 101, 673–693;
(c) Cornish-Bowden, A. J Biosci 2002, 27, 121–126.
21. (a) Olah, G. A.; Kuhn, S. J. Org Synth Coll 1973, 5, 480;
(b) Olah, G. A.; Kuhn, S. J. Org Synth Coll 1967, 47,
56; (c) Surya Prakash, G. K.; Mathew, T. Angew Chem,
Int 2010, 49, 1726.
International Journal of Chemical Kinetics DOI 10.1002/kin.21068