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172
Chapter – 4
Discussion
The Kinetics is a branch of science which deals with rate and
mechanism of chemical reaction .Literature survey reveals that the
kinetics of oxidation of organic compounds have been studied by
many researchers (48–52), survey also reveals that less work have
been done in inorganic reaction mechanism but very less work has
been carried out in the field of precipitation reactions.
For the present study, we have taken a definite amount of
solid sodium thiosulphate which is treated with hydrochloric acid
solution. The solution becomes turbid, due to the formation of
Sulphur particles. The reaction takes sufficient time; therefore it
can be successfully studied by using nephelometry. The principle
of nephelometry is based on scattering of light. The nephelometer
was calibrated using formazine solution. The rate of precipitation,
as is well established is governed by various factors such as
solubility product, rate of nucleation; type of salt formed, solvent
used, temperature, etc.
The reaction of sodium thiosulphate with hydrochloric acid
is a precipitation reaction, and stoichiometry of reaction is
Na2S2O3 + 2HCl 2NaCl + SO2 + H2O + S
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The probable mechanism of the precipitation reaction can be
given as follows:
(1) 2HCl 2H + 2Cl
O O O O H
Na S S Na Na S S Na
O O
OH
Na SO2 S Na
H OH2 Na SO2 S + Na Cl O
Na S S + NaCl + H2O
O Cl
NaCl + SO2 + S
Sodium thiosulphate and hydrochloric acid were allowed to
mix in definite proportion. The NTU reading at various time
intervals were recorded and it was observed that NTU reading
increases with the increasing concentration of sodium thiosulphate.
By using first order rate expression values of rate constant
were calculated. For the determination of rate constant, readings
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corresponding to initial concentration was taken as
[(NTU)∞ – (NTU)0] and reading corresponding to remaining
concentration of reactant was taken as [(NTU) ∞ – (NTU)t].
A graph of log a / a – x Vs time was plotted which passes
through origin, which shows first order reaction with respect to
concentration. The slope of line is used to calculate rate constant
values for different concentration of sodium thiosulphate.
The variation of rate constant with concentration of sodium
thiosulphate shows a linear dependent.
Table: 4.1
Variation of rate constant with [Na2S2O3]:
[HCl] (M)
[Na2S2O3]
1.343 x 10– 2M k Sec–1
[Na2S2O3]
2.686 x 10– 2M k Sec–1
[Na2S2O3]
4.029 x 10– 2M k Sec–1
0.1 2.3 x 10– 3 2.9 x 10– 3 14.5 x 10– 3
0.05 4.6 x 10– 3 6.2 x 10– 3 7.5 x 10– 3
0.025 6.3 x 10– 3 9.9 x 10– 3 31.5 x 10– 3
From these values it is clear that the rate of reaction
increases with the increasing concentration of the reactant. The rate
constant varies with concentration of [HCl] and [Na2S2O3]. The
values are represented in table no: 4.1.
The increase in rate of reaction with the increase in
concentration of reactants can be explain on the basis of collision
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theory of reaction rate which says that the number of collision
increases with the increasing concentration, it is so because the
number of molecules per unit volume is increasing thereby
increasing the rate of collision among the molecules which gives
the product, and hence the rate of precipitation reactions increases.
This reaction rate is found to be of the first order kinetics.
Brijesh pare (53) and others have studied rates for the miceller
catalyzed oxidation of Ethanol which shows dependence of rate on
surfactants concentration. Pardeep K. Sharma (54) have studied
oxidation of formic acid and oxalic acid and showed that the rate of
reaction increases with the increase in concentration of these
organic acids that are dependent on concentration of reactants.
Sarju Prasad (55) and others have studied conductometric titration of
yttrium nitrate solution at concentration 0.02M Potassium
Chromate. Tareev and Baev(56) on the basis of mathematical
equation, have shown that electrical conductance is a function of
concentration and particle size of dispersed phase. In the present
work, we have studied precipitation reaction by nephelometry
where turbidity unit reading increases with concentration.
Varsha Dhariwal(57) and others have calculated rate constant
values for oxidation of unsaturated acid and have given ‘K’ values
for different concentration of crotonic acid. For 0.1mol dm–3 as
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5.81 x 10–4 Sec–1.We have calculated ‘K’ values for 4.029 x 10–3 M
solution of Sodium thiosulphate as 1.6 x 10–3.
From the values given in table: 4.1 it has also been observed
that the rate constant increases with the decrease in concentration
of [HCl]. Zahid Amjad(58) have studied the effect of solution pH on
the crystal growth of Calcium Sulphate in the presence of polymer
by using seeded growth technique. Considerable attention has been
given to the various forms of Calcium Sulphate crystallizing from
aqueous solution. It is affected by temperature and stoichiometric
ratio of lattice ions. Sheehan and Nancollas(59) have examined the
effect of pH in the range of 6 – 9 on the growth of Calcium Oxalate
CaC2O4 in the presence of polyacrylate. It was found that the rate
of growth was independent of solution pH.
This reaction of Sodium thiosulphate and hydrochloric acid
has also been observed for the effect of salt KCl. It is observed that
with few exceptions the rate of reaction increases with increasing
concentration of salt KCl.
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Table: 4.2
Effect of salt [KCl] on rate constant:
[KCl] (M)
[HCl] = 0.1M [Na2S2O3]
1.343 x 10– 2M k Sec–1
[HCl] = 0.1M [Na2S2O3]
2.686 x 10– 2M k Sec–1
[HCl] = 0.1M [Na2S2O3]
4.029 x 10– 2M k Sec–1
0.0 2.3 x 10– 3 2.9 x 10– 3 1.45 x 10– 3
0.1 2.2 x 10– 3 4.3 x 10– 3 1.1 x 10– 3
0.2 8 x 10– 3 5.2 x 10– 3 1.12 x 10– 3
0.3 8.9 x 10– 3 8.4 x 10– 3 0.85 x 10– 3
Krishna Chandra (60) and others have studied stability of
lyophobic colloids. According to him, theoretical expression is
ordinarily not feasible because neither there is sufficient
knowledge(61) of a number of theoretical evolution in the relation
obtained by Verwey and Overbeek(62), nor the process of
coagulation is completely devoid of any specific effect of the ions
of added electrolyte.
The reaction has been studied at different temperatures i.e.
288 K, 298 K, 303 K and ‘K’ values at these temperatures were
found to be 3.6 x 10–3 Sec–1, 4.6 x 10–3 Sec–1 and 5 x 10–3 Sec–1. It
has been observed that the rate of precipitation increases with
increasing temperature, the reason for this can be given on the basis
of molecular theory of collision of reaction rate which says that as
the temperature is increased the particles get extra energy and they
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move faster. It means that the number of collision among the
molecules is increasing and when this collision is in a proper
orientation then the rate of formation of precipitate increases and
hence the rate is increased. When the reaction is studied at room
temperature less number of molecules passes the activation energy
barrier. This reaction of sodium thiosulphate and hydrochloric acid
is exothermic, which push some of the particles beyond their
activation energy. Because of this more number of the particles of
sodium thiosulphate becomes available to react with HCl. At
higher temperature more number of particles becomes activated
and process becomes very fast thereby increasing the rate of
reaction. B.P. Yadava(63) and others have studied that there is
increasing specific conductance and velocity of Al(OH)3 at higher
temperatures.
At different temperatures, rate constants were calculated, and
from the linear Arrhenius plot of log K Vs 1/ T, the computed
activation parameter for overall reaction were evaluated.
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Table: 4.3
Temperature Effect:
[Na2S2O3] = 2.686 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1
Log K
288 0.003472 3.6 x 10-3 -2.4437
298 0.003356 4.6 x 10-3 -2.33724
303 0.0033 5 x 10-3 -2.30103
From the values of rate constant activation energy of
reaction were calculated by using Arrhenius equation i.e.
k = Ae – E / RT
Free energy, ∆G* is calculated by using equation
∆G* = – RTlnK
Free energy values at different temperatures were calculated
and graph of ∆G* Vs Temperature was plotted which gives straight
line, slope of this line is used to calculate ∆S* values i.e. entropy
change, intercept of this line gives values of ∆H*.
Ea* = 23.51 x 102 KJ mol–1 ∆G* = 36.652 x 102 KJ mol–1
∆H* = 36.15 x 102 KJ mol–1 ∆S* = –288.859 KJ mol–1
For higher concentration, the effect of temperature is written as
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Table: 4.4
Temperature effect
[Na2S2O3] = 4.029 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 5x10-3 -2.30103
298 0.003356 5.2x10-3 -2.284
303 0.0033 10.5x10-3 -1.97881
Table: 4.5
Temperature effect
[Na2S2O3] = 5.372 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 7.5x10-3 -2.12494
298 0.003356 10.3x10-3 -1.98716
303 0.0033 17.7x10-3 -1.75203
The temperature dependence of rate constant is responsible
for temperature of rate of reaction. Arrhenius noticed that the
magnitude of temperature effect on reaction rate was too large to
be explained in terms of only a change in the translational energy
of the reactants. Thus, for a reaction to occur, it requires more than
just a collision between the reactants.
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Although the Arrhenius equation is used extensively to
determine the activation energies of chemical reactions, the plot of
lnk Vs 1 / T for some reactions is not linear, such non linear
behavior can now be justified, and many modern theories of
reaction rates predict that rate constant behaves like
K = a T m e –E / RT
Where a, E, and m are temperature dependent constants.
This reaction of sodium thiosulphate and hydrochloric acid
has also been observed for the effect of radiation of light. It is
noted that there is no effect of radiation on the rate of precipitation.
It indicates that precipitation is a thermal reaction and not a photo
chemical reaction.
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Cerium Chloride:
The reaction of Cerium Chloride with Oxalic acid is a
precipitation reaction. In the present study Cerium Chloride in
different concentration was taken and it is reacted with oxalic acid.
The stoichiometry of reaction is,
2Ce +2 + 2(COOH)2 2Ce(COO)4
The probable mechanism of the precipitation reaction can be
given as follows:
1) COOH COO + H COOH COOH + H 2) CeCl3 CeCl2 + HCl 3) COO COO + CeCl2 Ce + HCl COOH COO Cerium Chloride and Oxalic acid was allowed to mix in
definite proportion. The NTU readings at various time intervals
were recorded and it was observed that NTU reading increases with
the increasing concentration of Cerium Chloride.
By using first order rate expression, values of rate constant
were calculated. For the determination of rate constant, readings
corresponding to initial concentration was taken as
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[(NTU)∞ – (NTU)0] and reading corresponding to remaining
concentration of reactant was taken as [(NTU) ∞ – (NTU)t].
A graph of log a / a – x Vs time was plotted which passes
through origin, which shows first order reaction with respect to
concentration. The slope of line is used to calculate rate constant
values for different concentration of Cerium Chloride.
The variation of rate constant with concentration of Cerium
Chloride shows a linear dependent.
Table: 4.6
Variation of rate constant with [CeCl3]:
[(COOH) 2] (M)
[CeCl3]
1.352 x 10– 2M k Sec–1
[CeCl3]
2.704 x 10– 2M k Sec–1
[CeCl3]
4.056 x 10– 2M k Sec–1
0.1 8.9 x 10– 3 9.6 x 10– 3 3.84 x 10– 2
0.05 2.04 x 10– 2 5.041 x 10– 2 3.06 x 10– 1
0.025 9.2 x 10– 3 1.036 x 10– 2 2.03 x 10– 2
From these values it is clear that the rate of reaction
increases with the increasing concentration of reactant. The rate
constant varies with concentration of [(COOH)2] and [CeCl3]. The
values are represented in table no: 4.6.
The increase in rate of reaction with the increase in
concentration of reactants can be explained on the basis of collision
theory of reaction rate which says that the number of collision
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increases with the increasing concentration, it is so because the
number of molecules per unit volume is increasing there by
increasing the rate of collision among the molecules which gives
the product, and hence the rate of precipitation reactions increases,
this reaction rate is found to be of first order kinetics.
Jamil Ahmed (64) have studied precipitation of sodium acetate
trihydrate and observed that the rate of crystallization was found to
increase with the relative super saturation. S. A. Chimatadar(65) and
others have calculated initial rates of oxidation of ninhydrin by
quinolinium dichromate. For the concentration 0.4 x 103 mole dm–3
the initial rate is 3.31 mole dm–3 Sec–1. In our study when cerium
chloride concentration is 1.352 x 10–2 mole then initial rate is found
to be 2.3 moles Sec–1. He also have reported order of reaction for
different concentrations of ninhydrin as, when concentration is
4 x 10–4 mole dm–3 order is one and when concentration is 8 x 10–2
mole dm–3 order is 0.67. Whereas for Cerium Chloride reaction we
have calculated order it is given in table: 4.6. Gosta Benglsson(66)
and others have reported determination of rate law by method of
initial rates using solution. Initially, containing comparable
concentration of ion (III) and ion (II) together with access of
hydroxyl ammonium. Under these conditions, reaction becomes
sufficiently slow for values of υ0 and shows that initial rate
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depends on concentration of Fe (III) and Fe (II). S. N. Dindi(67) and
others have reported oxidation of Tellurium by Cerium and showed
it as a first order reaction. Different concentrations of Cerium (III)
also give first order kinetics.
The rate of reaction increases with the decreasing
concentration of oxalic acid, showing dependence of rate on pH of
solution. I.P. Saraswat(68) and others studied solubility of
ammonium manganese phosphate., and observed that for lower
value of pH of solution the concentration of H+ ions will be higher
and hence PO4– –ions would tend to decrease due to association,
phosphoric acid being a weak acid. Thus, both the reactions would
be forced towards the right and hence higher solubility of
compound is expected with decrease in pH. By adding free Mn++
ions in the solution the solubility of ammonium manganese
phosphate decreases as is to be expected, because with the addition
of free Mn++ ions the reaction should be forced towards left.
The reaction has been studied at different temperatures i.e.
288 K, 298 K, 303 K and ‘K’ values at these temperatures were
found to be 3.6 x 10–3 Sec–1, 4.6 x 10–3 Sec–1 and 5 x 10–3 Sec–1. It
has been observed that the rate of precipitation increases with the
increasing temperature, the reason for this can be given on the basis
of molecular theory of collision of reaction rate which says that as
186
the temperature is increased the particles get extra energy and they
move faster. It means that the number of collision among the
molecules is increasing and when this collision is in a proper
orientation then the rate of formation of precipitate increases and
hence the rate is increased. When the reaction is studied at room
temperature less number of molecules passes the activation energy
barrier. This reaction of Cerium Chloride and Oxalic acid is
exothermic, which push some of the particles beyond their
activation energy. Because of this more number of the particles of
Cerium Chloride becomes available to react with Oxalic acid. At
higher temperature more number of particles becomes activated
and process becomes very fast thereby increasing the rate of
reaction. Rupal Kumbhat(69) and others have carried out oxidation
of formic acid by quinolinium bromochromate . He says that
reaction rate increases with the increase in organic acid and also
increases with the increase in temperature. He had calculated rate
constant at 288, 298 and 308 K for formic acid as 18.4 dm–3 mole–1
S–1, 38.1 dm–3 mole–1 S–1, 75.7 dm–3 mole–1 S–1 respectively. Rate
constant were calculated at different temperatures (table 4.7)
At different temperatures, rate constant were calculated, and
from the linear Arrhenius plot of log K Vs 1 / T, the computed
activation parameter for overall reaction were evaluated.
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Table: 4.7
Temperature Effect
[CeCl3] = 1.352 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1
Log K
288 0.003472 1.55 x10-2 -1.809
298 0.003356 1.58 x 10-2 -1.801
303 0.0033 1.7 x10-2 -1.769
From the values of rate constant, activation energies of
reaction were calculated by using Arrhenius equation i.e.
k = Ae – E / RT
Free energy, ∆G* is calculated by using equation
R.L. Yadav and Verma(70) have reported Ceric–Cerium
oxidation of Catchechol and said that Arrhenious law was found to
be valid.
∆G* = – RTlnK
Free energy values at different temperatures were calculated
and graph of ∆G* Vs Temperature was plotted which gives straight
line, slope of this line is used to calculate ∆S* values i.e. entropy
change, intercept of this line gives values of ∆H*.
Ea* = 34.82 x 102 KJ mol–1 ∆G* = 30.51 x 102 KJ mol–1
∆H* = 29.75 x 102 KJ mol–1 ∆S* = –30.2 KJ mol–1
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M. Bala Krishnan (71) and others have calculated ∆G*, ∆S*
and ∆H* at different temperatures as 92.66 kj mole–1,–158.56 J
mole–1, K–1, 45.25 kj mole–1, respectively. Sheila Srivastava(72)and
others have reported the reaction is first order with respect to RIII
and the values of rate constant were calculated. The graph of log K
Vs 1 / T was plotted to calculate the value of ∆G*, ∆S*, ∆H* and it
is given as 56 kj mole–1 and 58 kj mole–1, 69 kj mole–1. Samir
Kumar(73) and others have studied coagulation and electrical
properties of Ceric oxide hydrosols and its conductance were
measured which was found to be 1.04 x 10–4 Ohms at 250C.
Temperature significantly affects the flocculation (74), higher
temperature adversely affect the flocculation as evidenced by lower
settling rates at higher temperature. Variation in temperature alters
intrinsic viscosity of the polymer due to which settling rate also
changes.
Table: 4.8
Temperature effect
[CeCl3] = 2.704 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 3.40 x10–2 -1.46852
298 0.003356 3.47 x10–2 -1.45967
303 0.0033 3.8 x10–2 -1.42022
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Table: 4.9
Temperature effect
[CeCl3] = 4.056 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 2.83 x10–2 -1.54821
298 0.003356 1.6 x10–1 -0.79588
303 0.0033 1.65 x10–1 -0.78252
The temperature dependence of rate constant is responsible
for the temperature of rate of reaction. Arrhenius noticed that the
magnitude of temperature affect on reaction rate was too large to
be explained in terms of only a change in the translational energy
of the reactants. Thus for a reaction to occur, it requires more than
just a collision between reactants.
Although the Arrhenius equation is used extensively to
determine the activation energies of chemical reactions, the plot of
lnk Vs 1 / T for some reactions is not linear, such non linear
behavior can now be justified, and many modern theories of
reaction rates predict that rate constant behaves like
K = a T m e –E / RT
Where a, E, and m are temperature dependent constants.
190
This reaction of Cerium Chloride and Oxalic acid has also
been observed for the effect of radiation of light. It is noted that
there is no effect of radiation on the rate of precipitation. It
indicates that precipitation is a thermal reaction and not a photo
chemical reaction.
This reaction of Cerium Chloride and Oxalic acid has also
been observed for the effect of salt KCl. It is observed that with
few exceptions the rate of reaction increases with increasing
concentration of salt KCl.
Table: 4.10
Effect of salt [KCl] on rate constant
[KCl] (M)
[(COOH) 2] = 0.1M [CeCl3]
1.352 x 10– 2M k Sec–1
[(COOH) 2] = 0.1M [CeCl3]
2.704 x 10– 2M k Sec–1
[(COOH) 2] = 0.1M [CeCl3]
4.056 x 10– 2M k Sec–1
0.0 3.84 x 10– 2 8.9 x 10– 3 9.6 x 10– 3
0.1 1.28 x 10– 2 9.67 x 10– 3 1.26 x 10– 2
0.2 1.31 x 10– 2 1.70 x 10– 2 1.28 x 10– 2
0.3 2.99 x 10– 2 4.49 x 10– 2 1.47 x 10– 2
Ammonium molybdate(75) solution was coagulated by
thorium nitrate and mixed with non electrolyte this mixture was
studied by using turbidimeter. And variation in turbidity with time
was followed by reading the absorbance at intervals. He reported
that turbidity is directly proportional to the decreasing percentage
191
of transmittance. He plotted the graph of decreasing percentage
transmit tens Vs Time and the graph showed rise in turbidity with
time. By these studies he reported that the mechanism of slow
coagulation of molybdenum blue solution by thorium nitrate is the
same in presence and absence of electrolyte which indicates that
there is no effect of electrolyte on coagulation process. Mukhtar
Singh(76) and others have studied coagulation of nickel hydrous
oxide solution. He has plotted 1 / c – a against Time and it is a
straight line graph.
This reaction of Cerium Chloride and Oxalic acid has also
been observed for the effect of radiation of light. It is noted that
there is no effect of radiation on the rate of precipitation. It
indicates that precipitation is a thermal reaction and not a photo
chemical reaction.
192
Strontium Chloride:
The reaction of Strontium chloride with Sulphuric acid is a
precipitation reaction. In the present study I have taken Strontium
chloride in different concentration was taken and it was reacted
with Sulphuric acid. The stochiometry of reaction,
Sr+2 + SO4-2 SrSO4
The probable mechanism of the precipitation reaction can be
given as follows:
1) H2SO4 H + HSO4 2) SrCl2 + H HCl + SrCl 3) SrCl + HSO4 HCl + SrSO4
Strontium Chloride and Sulphuric acid was allowed to mix
in definite proportion. The NTU readings at various time intervals
were recorded and it was observed that NTU reading increases with
the increasing concentration of Strontium Chloride.
By using first order rate expression, values of rate constant
were calculated. For the determination of rate constant, readings
corresponding to initial concentration was taken as
[(NTU)∞ – (NTU)0] and reading corresponding to remaining
concentration of reactant was taken as [(NTU) ∞ – (NTU)t].
A graph of log a / a – x Vs time was plotted which passes
through origin, which shows first order reaction with respect to
193
concentration. The slope of line is used to calculate rate constant
values for different concentration of Strontium Chloride.
The variation of rate constant with concentration of
Strontium Chloride shows a linear dependent.
Table: 4.11
Variation of rate constant with [SrCl2]
[H 2SO4] (M)
[SrCl 2]
2.1028 x 10– 2M k Sec–1
[SrCl 2]
4.2056 x 10– 2M k Sec–1
[SrCl 2]
6.3085 x 10– 2M k Sec–1
0.1 9.5 x 10– 3 12.3 x 10– 3 24.4 x 10– 3
0.05 7.5 x 10– 3 12.8 x 10– 3 14.0 x 10– 3
0.025 5.6 x 10– 3 8.9 x 10– 3 10.1 x 10– 3
From these values it is clear that the rate of reaction
increases with the increasing concentration of the reactant. The rate
constant varies with concentration of [H2SO4] and [SrCl2]. The
values are represented in table no: 4.11.
The increase in rate of reaction with the increase in
concentration of reactants can be explained on the basis of collision
theory of reaction rate which says that the number of collision
increases with the increasing concentration, it is so because the
number of molecules per unit volume is increasing thereby
increasing the rate of collision among the molecules which gives
194
the product, and hence the rate of precipitation reactions increases,
this reaction rate is found to be of first order kinetics.
Octakar Suhnel(77) have studied precipitation of Strontium
Sulphate. In this work, they have studied precipitation, growth and
nucleation of Strontium Sulphate crystals in super saturated
solution. He showed that this reaction is a second order kinetics
with respect to concentration and temperature. The difference
between the study carried out by Octakar and present work is that,
when Sulphuric acid is allowed to react with Strontium chloride
precipitation occurs which follows the first order kinetics in present
experimental condition .Dr. Octakar studied the precipitation of
Strontium Sulphate from the aqueous solution without seeding
under super saturation condition.
F. Johns (78) have studied the mechanism of nitrolotriacetic
acid (NTA) in which they observed interaction of Barium Sulphate
with (NTA) which enhances the precipitation of Barium Sulphate.
The reaction of Strontium Chloride and Sulphuric acid has
also been studied for effect of salt KCl. It is observed that with few
exceptions the rate of reaction increases with increasing
concentration of salt KCl.
195
Table: 4.12
Effect of salt [KCl] on rate constant
[KCl] (M)
[H 2SO4] = 0.1M [SrCl 2]
4.2056 x 10– 2M k Sec–1
[H 2SO4] = 0.1M [SrCl 2]
6.3084 x 10– 2M k Sec–1
[H 2SO4] = 0.1M [SrCl 2]
8.4111 x 10– 2M k Sec–1
0.0 12.3 x 10– 3 24.4 x 10– 3 20.2 x 10– 3
0.1 1.22 x 10– 2 9.1 x 10– 3 3.2 x 10– 3
0.2 8.2 x 10– 3 11.1 x 10– 3 10.1 x 10– 3
0.3 8.21 x 10– 3 7.1 x 10– 3 3.15 x 10– 2
The reaction has been studied at different temperature i.e.
288 K, 298 K, 303 K and ‘K’ values at these temperatures were
found to be 2.9 x 10–3 Sec–1, 3.9 x 10–3 Sec–1 and 11.5 x 10–3 Sec–1.
It has been observed that the rate of precipitation increases with the
increasing temperature, the reason for this can be given on the basis
of molecular theory of collision of reaction rate which says that as
the temperature is increased the particles get extra energy and they
move faster. It means that the number of collision among the
molecules is increasing and when this collision is in a proper
orientation then the rate of formation of precipitate increases and
hence the rate is increased. When the reaction is studied at room
temperature less number of molecules passes the activation energy
barrier. This reaction of Strontium Chloride and Sulphuric acid is
exothermic, which push some of the particles beyond their
196
activation energy. Because of this more number of the particles of
Strontium Chloride becomes available to react with HCl. At higher
temperature, more number of particles becomes activated and
process becomes very fast thereby increasing rate of reaction.
At different temperatures rate constant were calculated, and
from the linear Arrhenius plot of log K Vs 1/ T, the computed
activation parameter for overall reaction were evaluated.
Table: 4.13
Temperature Effect
[SrCl2] = 2.1028 x 10–2 M
Temp K 1 / T Rate Constant
k Sec–1 Log K
288 0.003472 2.9 x 10-3 -2.5376
298 0.003356 3.9 x 10-3 -2.40894
303 0.0033 11.5 x 10-3 -1.9393
From the values of rate constant activation energy of
reaction were calculated by using Arrhenius equation i.e.
k = Ae – E / RT
Free energy, ∆G* is calculated by using equation
∆G* = – RTlnK
Free energy values at different temperature were calculated
and graph of ∆G* Vs Temperature was plotted which gives straight
197
line, slope of this line is used to calculate ∆S* values i.e. entropy
change, intercept of this line gives values of ∆H*.
Ea* = 30.24 x 102 KJ mol–1 ∆G* = 32.10 x 102 KJ mol–1
∆H* = 32.40 x 102 KJ mol–1 ∆S* = –86.66 KJ mol–1
Vidhya Sagar(79) and others have carried out conductometric
investigation of thorium soaps at different temperatures .They have
studied butyrate soap at 400C, 500C and 600C free energy values at
these temperature were calculated and it is given as 40.032KJ mol-1
41.445 KJ mol-1, 43.004KJ mol-1.
D.R. Srivastava(80) and others have studied specific
interaction of m-dinitrobenzen with naphthalene and they have
calculated enthalpy at different temperatures i.e. 25.40C, 30.20C,
35.60C these values are 1760 Cal, 1746 Cal, 1652 Cal respectively
For higher concentration, the effect of temperature is written as:
198
Table: 4.14
Temperature effect
[SrCl2] = 4.2056 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 3.91 x10-3 -2.40782
298 0.003356 4.8 x10-3 -2.31876
303 0.0033 13.3 x10-3 -1.87615
Table: 4.15
Temperature effect
[SrCl2] = 6.3084 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 10.8 x10-3 -1.96658
298 0.003356 12.2 x10-3 -1.91364
303 0.0033 20.2 x10-3 -1.69465
The temperature dependence of rate constant is responsible
for temperature of rate of reaction. Arrhenius noticed that the
magnitude of temperature effect on reaction rate was too large to
be explained in terms of only a change in the translational energy
of the reactants. Thus for a reaction to occur, it requires more than
just a collision between reactants.
199
Although the Arrhenius equation is used extensively to
determine the activation energies of chemical reactions, the plot of
lnk Vs 1 / T for some reactions is not linear, such non linear
behavior can now be justified, and many modern theories of
reactions rates predict that rate constant behaves like
K = a T m e –E / RT
Where a, E, and m are temperature dependent constants.
This reaction of Strontium Chloride and Sulphuric acid has
also been observed for the effect of radiation of light. It is noted
that there is no effect of radiation on the rate of precipitation. It
indicates that precipitation is a thermal reaction and not a photo
chemical reaction.
200
Silver Nitrate:
The reaction of Silver Nitrate with hydrochloric acid is a
precipitation reaction. In the present study Silver Nitrate was taken
in different concentrations and it was reacted with hydrochloric
acid which gives precipitate. The stochiometry of reaction is,
Ag+ + Cl– AgCl
The probable mechanism of the precipitation reaction can be
given as follows:
1) HCl H + Cl 2) AgNO3 + H HNO3 + Ag 3) Ag + Cl AgCl
Silver Nitrate and Hydrochloric acid was allowed to mix in a
definite proportion. The NTU readings at various time intervals
were recorded and it was observed that NTU reading increases with
the increasing concentration of Silver Nitrate.
By using first order rate expression, values of rate constant
were calculated. For the determination of rate constant, readings
corresponding to initial concentration was taken as
[(NTU)∞ – (NTU)0] and reading corresponding to remaining
concentration of reactant was taken as [(NTU) ∞ – (NTU)t].
201
A graph of log a / a – x Vs time was plotted which passes
through origin, which shows first order reaction with respect to
concentration. The slope of line is used to calculate rate constant
values for different concentration of Silver Nitrate.
The variation of rate constant with concentration of Silver
Nitrate shows a linear dependent.
Table: 4.16
Variation of rate constant with [AgNO3]:
[HCl] (M)
[AgNO3]
1.9623 x 10– 2M k Sec–1
[AgNO3]
3.9246 x 10– 2M k Sec–1
[AgNO3]
5.8869 x 10– 2M k Sec–1
0.1 0.5 x 10– 2 1.45 x 10– 2 2.51 x 10– 2
0.05 2.2 x 10– 2 4.07 x 10– 2 9.9 x 10– 2
0.025 2.9 x 10– 3 6.8 x 10– 3 8.7 x 10– 3
From these values it is clear that the rate of reaction
increases with the increasing concentration of the reactant. The rate
constant varies with concentration of [HCl] and [AgNO3]. The
values are represented in table no: 4.16.
M. K. Mishra(81) and others have studied oxidation of
polyvinyl alcohols by benzyl triethyl ammonium chlorochromate
and have calculated K value for different concentration of alcohols
202
as 0.1 mol dm–3, 0.2, 0.4, at 298 K as 13.7 S–1, 27.65S–1, 55.1 S–1,
respectively.
The increase in rate of reaction with the increase in
concentration of reactants can be explained on the basis of collision
theory of reaction rate which says that the number of collision
increases with the increasing concentration, it is so because the
number of molecules per unit volume is increasing thereby
increasing the rate of collision among the molecules which gives
the product, and hence the rate of precipitation reactions increases,
this reaction rate is found to be of first order kinetics.
Harkanwal Singh (82) and others have studied precipitation
reaction of barium chloride to give Barium molibdate by using
simple precipitation study and weight of this precipitate was
determined. During coagulation of Silver Chloride (83) with the
electrolytes of varying valancies observed that the adsorption was
high in stable region and vice versa. Dr. Banargea(84) have carried
out kinetics study of inorganic reactions. He used
spectrophotometric method for the determination of order of
reaction. Measurement of rate at different acid concentration
indicates that following expression holds goods.
Rate = K abs [Complex] [H+]
203
This reaction of Silver Nitrate and Hydrochloric acid has
also been observed for the effect of salt KCl. It is observed that
with few exceptions the rate of reaction increases with increasing
concentration of salt KCl. Flocculation studies of ferric vanadate
was carried out by Krishna Raina(85) and others. They observed
these by extinction measurement, and stated that the changes take
place with step wise addition of electrolyte. In the beginning there
is no change in extinction.
In the present study effect of salt on the precipitation
reaction is given table no: 4.17.
Table: 4.17
Effect of salt [KCl] on rate constant
[KCl] (M)
[HCl] = 0.1M [AgNO3]
1.9623 x 10– 2M k Sec–1
[HCl] = 0.1M [AgNO3]
3.9246 x 10– 2M k Sec–1
[HCl] = 0.1M [AgNO3]
5.8869 x 10– 2M k Sec–1
0.0 0.5 x 10– 2 1.45 x 10– 2 2.51 x 10– 2
0.1 1.2 x 10– 3 2.5 x 10– 3 2.4 x 10– 3
0.2 0.92 x 10– 3 6.9 x 10– 3 4.8 x 10– 3
0.3 3.2 x 10– 3 2.3 x 10– 3 2.5 x 10– 3
At different temperatures rate constant were calculated, and
from the linear Arrhenius plot of log K Vs 1/ T, the computed
activation parameter for overall reaction were evaluated.
204
Table: 4.18
Temperature Effect
[AgNO3] = 1.9623 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1
Log K
288 0.003472 7.5 x 10-3 -2.1249387
298 0.003356 8.5 x 10-3 -2.0705811
303 0.0033 2.2 x 10-2 -1.6575773
From the values of rate constant activation energies of
reaction were calculated by using Arrhenius equation i.e.
k = Ae – E / RT
Free energy, ∆G* is calculated by using equation
∆G* = – RTlnK
Free energy values at different temperatures were calculated
and graph of ∆G* Vs Temperature was plotted which gives straight
line, slope of this line is used to calculate ∆S* values i.e. entropy
change, intercept of this line gives values of ∆H*.
Ea* = 27.62 x 102 KJ mol–1 ∆G* = 39.29 x 102 KJ mol–1
∆H* =35.50 x 102 KJ mol–1 ∆S* = –28.93 KJ mol–1
Kavita Chawan(86) and others have carried out studies for
aliphatic aldehyde by benzyltriethylammonium chlorochromate
205
and calculated values for ∆G*, ∆S*, ∆H*, as 92.6 0.6, KJ mol–1
– 94 2 J mol–1 K–1, 64.6 0.7 KJ mol–1.
For higher concentration, the effect of temperature is written as:
Table: 4.19
Temperature effect
[AgNO3] = 3.9246 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 8.5 x 10-3 -2.07058
298 0.003356 8.52 x 10-3 -2.06956
303 0.0033 15.6 x 10-3 -1.80688
Table: 4.20
Temperature effect
[AgNO3] = 5.8869 x 10–2 M
Temp K 1 / T Rate Constant k Sec–1 Log K
288 0.003472 3.2 x10-3 -2.49485
298 0.003356 1.4 x10-2 -1.85387
303 0.0033 1.65 x10-2 -1.78252
The temperature dependence of rate constant is responsible
for temperature of rate of reaction. Arrhenius noticed that the
magnitude of temperature effect on reaction rate was too large to
206
be explained in terms of only a change in the translational energy
of the reactants. Thus, for a reaction to occur, it requires more than
just a collision between reactants.
Although the Arrhenius equation is used extensively to
determine the activation energies of chemical reactions, the plot of
lnk Vs 1 / T for some reactions is not linear, such non linear
behavior can now be justified, and many modern theories of
reactions rates predict that rate constant behaves like
K = a T m e –E / RT
Where a, E, and m are temperature dependent constants.
This reaction of Silver Nitrate and Hydrochloric acid has
also been observed for the effect of radiation of light. It is noted
that there is no effect of radiation on the rate of precipitation. It
indicates that precipitation is a thermal reaction and not a photo
chemical reaction.
207
Induced Reaction:
All the permanganate oxidations are usually complicated
because of different oxidation state of Mn (+7, +6, +5, +4, +3, +2),
that can participate in the reaction. In induced reaction, each
species has its own induction factor (IF), theoretically it is possible
to calculate it by using an equation
Where Ai --- oxidation state of intermediate of the actor.
Af ----final oxidation state of actor.
A0 ----initial oxidation state of actor.
By using above equation we can predict the probable
species, which can induce the reaction. The species with their
theoretical induction factors are given in the following table.
Table: 4.21
Manganese species of varied oxidation states with their induction
factors.
I. F
Mn+6 4.00
Mn+5 1.50
Mn+4 0.66
Mn+3 0.25
Ai – Af IF = A0 – Ai
208
It is observed that if the oxidation of Ce+3 is carried out in
the presence of oxalic acid, permanganate is consumed in access of
the amount required by Ce+3. Thus the over consumption of
permanganate which was observed here is due to an induced
oxidation of oxalic acid involving and produced in the
permanganate – Ce+3 reaction.
The induced reaction can be characterized by the induction
factor i.e. the ratio of oxidation is equivalently consumed by the
equivalents of the acceptor and the inductor, the induction factor
determine under varied experimental conditions are given in
different table(4.21).
The induction factor approaches to the value of 0.25 with the
increasing concentration of oxalic acid and it decreases rapidly in
other cases where no induced oxidation reaction was observed.
In the absence of oxalic acid, Ce+3 reduces permanganate ion
very rapidly to a mixture of Mn+3 and Mn+4, the presence of oxalic
acid in the reaction system will reduced both Mn+3 and Mn+4 to
Mn+2 as no evidence regarding the formation of Mn+4 was observe.
It is clear that whatever amount of Mn+4 is formed through either
this disproportionate. If Mn+3 are considered to be the end product
in acid condition employed in the reaction, these are not sufficient
209
to keep Mn+3 intact in the solution. Thus, as soon as it is formed, it
disproportionate to Mn+2.
The induced oxidation reactions have also been studied in
the presence of acceptor Sn+2 and Cu+2. It has been observed that
induction factor reaches to the maximum value of 0.25 which
indicates the presence of Mn+3 species in the reaction. The Mn+3 in
the system will not be stabilizing. As soon as it is found it will
immediately disproportionate to Mn+2 and Mn+4. However, it is
experimentally observed that with the solution after titration it
turns to yellow orange colour precipitate.
Effect of Salt:
The effect of salt such as sodium nitrate and Zinc Sulphate
was studied to understand the role in the characterization of
induced reaction. The induction factor was found to be 0.073
incase of sodium nitrate and 0.12 in case of Zinc Sulphate under
similar condition increasing induction factor shows following order
IF (NaNO2) < IF (ZnSo4)
When Cerium is used as an inductor then the induction
factor without the presence of the salt was found to be 0.08 and
when salt is used in the same oxidation reaction then the induction
factor is found to be 0.73 for NaNo2 and 0.12 for ZnSo4 which
210
clearly indicates that there is not much effect on induced oxidation
in the presence of salt.
211
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