Effect of Temperature on the Reaction Rate
Christy Joy A. Retanal, Vincent S. Dagala, Carlo Manuel A. Banquerigo
Chemical Engineering Department,
Xavier University - Ateneo de Cagayan
Corrales Avenue, Cagayan de Oro, Philippines
Abstract
The experiment aims to determine the activation energy
of the reaction, to determine the specific rate constant for
the reaction and lastly, to discuss the effect of
temperature on the reaction rate. On two Erlenmeyer
flasks, 100 mL HCl and 100 mL C4H8O2 were placed,
and then placed on a constant water bath at the selected
temperature (7.0, 28.9, 55⁰C) and allowed to attain
thermal equilibrium for about 5 minutes. After which, 5
mL of the ester was placed on the acid and the time for
mixing was recorded. For a certain temperature, 5 mL of
the reaction mixture was added to 75 mL iced H2O (For
different temperatures, different sampling schemes was
followed to fully observe the effects of temperature
change on the rate), time for mixing was recorded when
approximately half of the reaction mixture has flowed
down the pipette. Titration was done thereafter. It can be
seen from this experiment that rate of reaction is
concentration dependent while rate constant is not. Also,
that the rate constant for the hydrolysis of ethyl acetate
with sodium hydroxide using HCl as a catalyst is
approximately 0.00204min-1cm-3. The activation
energy of the reaction is 3497.58 J/mol which is
positive therefore signifies that the reaction rate
increases when temperature increases.
Keywords: specific rate constant, activation energy,
ionic reaction, reaction rate
I.
INTRODUCTION
The specific rate constant is the proportionality
constant in the equation that expresses the relationship
between the rate of a chemical reaction and the
concentrations of the reacting substances.[1] In other
words, at a given temperature, rate is equal to the rate
constant of reaction when concentration of the reactant
in unity. The expression of this relationship is;
d lnK E¿
=
dt
RT
Where K is the constant,
E¿
is the activation
energy, R is the ideal gas constant and T is the absolute
temperature. Assuming that the activation energy is
independent of temperature;
(T 2−T 1)
K2
Δ E¿
log
=
+
K 1 2.303 RT
T 2T 1
Collisions only result in a reaction if the
particles collide with enough energy to get the reaction
started. This minimum energy required is called the
activation energy for the reaction. Only those particles
represented by the area to the right of the activation
energy will react when they collide. The great majority
doesn’t have enough energy, and will simply bounce
apart.
To speed up the reaction, you need to increase
the number of the very energetic particles - those with
energies equal to or greater than the activation energy.
Increasing the temperature has exactly that effect - it
changes the shape of the graph.[2]
In summary, Increasing the temperature
increases reaction rates because of the disproportionately
large increase in the number of high energy collisions. It
is only these collisions (possessing at least the activation
energy for the reaction) which result in a reaction.
Experimentally, it is observed that rate of a
reaction increases with rise of temperature. This is
because as you increase the temperature, the kinetic
energy of the reactants increase, allowing for more
collisions between the molecules. This, therefore, allows
for products to be formed faster. [3]In general rate of a
reaction becomes double on rise of
(usually
).
Ionic Reactions are in most cases not adapted to
rate constant studies because of their high reaction rates.
For this experiment the hydrogen ion catalyzed
hydrolysis of ethyl acetate has been selected. Therefore,
the kinetics of the reaction can be studied by taking a
known quantity of ethyl acetate and mixing it with a
relatively large quantity of HCl. An aliquot of the
reaction mixture is withdrawn at different intervals of
time and titrated against a standard alkali. Obviously, as
the reaction proceeds, the value of alkali required to
neutralize the acid (HCl present as catalyst + CH3COOH
produced by hydrolysis of the ester) progressively
increases. [4]
II. For the intermediate temperature determination, the
first sample was taken at the end of 5 minutes. After that,
samples were taken in every 15 minutes in the span of 2
hours.
III. For the highest temperature determination, the first
sample was taken at the end of 5 five minutes. The next
samples were took in a 10 minute intervals for 30
minutes and at 30 minute intervals in a span of 2 hours.
These samples were drained into an Erlenmeyer
flask containing 75 mL of ice water. The time of
sampling when approximately half of the sample has
flowed from the sampling pipet was recorded. The
sample was then titrated. The volume of the titrant was
recorded.
III.
As a rough approximation, for many reactions
happening at around room temperature, the rate of
reaction doubles for every 10°C rise in temperature.
However, this is not true in all reactions and however,
increasing the temperature will not always increase the
rate of the reaction. If the temperature of a reaction were
to reach a certain point where the reactant will begin to
degrade, it will decrease the rate of the reaction. In this
experiment the effect of temperature on the reaction rate
is determined.
II.
RESULTS AND DISCUSSION
Reaction rate does depend on the rate constant of the
reactant, the concentration and the order of the reaction.
In this experiment, the rate constant was determined first
using the concentrations at a specific time and the use of
the equation:
2.303logc = kt + C
Eq. 1
which can also be represented as
(V o−V )
log (V t−V )
=
ktV
2.303
Vo
+ log V t
EXPERIMENTAL SECTION
The reaction was carried out at three different
set-ups; temperature on ice bath, room temperature and
high temperature. An Erlenmeyer flask with 100 mL of 1
N HCl acid and a flask containing ethyl acetate were
placed on a constant water bath at the selected
temperature. Both were allowed to settle in equilibrium
for 5 min. 5 mL of the ester was pipetted into the flask of
the acid. The time of mixing was recorded. 5 mL of the
reaction mixture was removed for analysis according to
the time scheduled given.
I. For the lowest temperature determinations, the first
sample was taken after 10 min. After that, the sample
was taken in every 25 minutes in the span of 2 hours.
Base on equations, we can generate a linear relationship
between the concentration of esther and the rate
constant. The c was first calculated which is the
concentration of ester at time t.
Table 1. Concentrations of Ester at 7.0⁰C
Time
(min)
10
35
60
85
110
135
mL NaOH
15.61
17.01
18.45
19.42
20.42
21.37
c=(Vo-V)/
(Vt-V)
1.00
1.32
1.97
2.95
6.06
-
logc
0
0.12
0.29
0.47
0.78
-
Plotting the values of log c as the y and time as the x, we
then make linear graph as shown in figure 1.
Figure 1. Graph of logc vs t at 7.0⁰C
(min)
5
15
25
35
65
(Vt-V)
1.00
1.31
1.79
4.88
-
17.40
18.60
19.64
21.43
22.47
0
0.12
0.25
0.69
-
0.8
Figure 3. Graph of logc vs t at 55.0⁰C
0.6
log c
f(x) = 0.02x - 0.18
0.4
0.2
0
0
10
20
30
0.8
Linear ()
0.6
Linear ()
0.4
40
log c
f(x) = 0.02x - 0.18
0.2
Time
Linear ()
Linear ()
0
0
The slope obtained from the graph is 0.022. Using this
slope we calculated for k and the value of k is min -1cm-3.
10
20
30
40
Time
Table 2. Concentrations of Ester at 28.9⁰C
Time
(min)
5
20
35
50
65
mL NaOH
16.29
16.47
16.61
16.80
17.06
c=(Vo-V)/
(Vt-V)
1.00
1.31
1.71
2.96
-
logc
0
0.12
0.23
0.47
-
Figure 2. Graph of logc vs t at 28.9⁰C
0.8
0.6
log c
f(x) = 0.02x - 0.18
0.4
0.2
Linear ()
Linear ()
0
0
10
20
30
40
Time
mL NaOH
It can be seen in the results obtained that the rate
constant of the hydrolysis of the ethyl acetate with HCl
as a catalyst was increasing on the three different
temperatures. Rate constant does depend on temperature.
Rates of most reactions are very sensitive to
temperature. Most increase rapidly with increasing
temperature. [5] This result follows the principle of
collision theory; if a substance is heated, the
particles move faster and so collide more frequently,
thus, speeding up the rate of reaction [6]. On
calculating the activation energy, the values of K
and temperature were used to plot a graph using
Arrhenius equation:
ln K  
Table 3. Concentrations of Ester at 55.0⁰C
Time
Table 4. Value of K at 7.0, 28.9 and 55.0⁰C
Temperature
Rate Constant
7.0⁰C
0.00204
28.9⁰C
0.00207
55.0⁰C
0.00254
c=(Vo-V)/
logc
Ea
 ln A
RT
Activation energy is the minimum energy which must be
available to a chemical system with potential reactants to
result in a chemical reaction. [7]
was not exact so the ratio of volume between the
reaction mixture and ice water is not accurate.
Figure 4. Graph of lnK vs 1/T
-5.85
-5.9 0
0
0
0
0
0
V.
0
-5.95
-6
-6.05
f(x) = - 420.66x - 4.72
Linear ()
-6.1
-6.15
CONCLUSION
It can be seen from this experiment that rate of reaction
is concentration dependent while rate constant is not.
Also, that the rate constant for the hydrolysis of ethyl
acetate with sodium hydroxide using HCl as a catalyst is
approximately 0.00204min-1cm-3. The activation
energy of the reaction is 3497.58 J/mol which is
positive therefore signifies that the reaction rate
increases when temperature increases.
-6.2
-6.25
REFERENCES
Using the slope of the graph which equal to
1
T
−Ea
log ¿=
R
slope ¿
k vs
the obtained activation energy was 3497.58 J/mol.
The determined activation energy was positive
therefore, the reaction rate is proportional to the
temperature. Also, the reaction needs 3497.58 J/mol
of energy in order for it to proceed into reaction.
Activation energy can be thought of as the height of
the potential barrier (sometimes called the energy
barrier) separating two minima of potential energy
(of the reactants and products of a reaction). For a
chemical reaction to proceed at a reasonable rate,
there should exist an appreciable number of
molecules with translational energy equal to or
greater than the activation energy. [7]
IV.
ERROR ANALYSIS
The errors that contributed in the inaccuracy in the
results on the experiment were: parallax error , delayed
in transferring the reaction mixture and the inexact
volume of reaction mixture pipetted. The volume of the
titrant recorded was not exact due to some samples that
were over titrated. The reaction mixture was not
transferred right-away and added to the volume of ice
water so there was a slightly change in temperature for
the mixture. The pipetted 5 mL of the reaction mixture
[1]http://www.britannica.com/science/reactionrate#ref274425
[2]http://www.chemguide.co.uk/physical/basicrates/temp
erature.html
[3]http://chemwiki.ucdavis.edu/Core/Physical_Chemistr
y/Kinetics/Reaction_Rates/Reaction_Rate#Temperature_
Dependence
[4]http://www.scihub.org/AJSIR/PDF/2015/1/AJSIR-61-1-4.pdf
[5]http://bouman.chem.georgetown.edu/S02/lect4/lect4.h
tm
[6]http://www.chemguide.co.uk/physical/basicrates/
temperature.html [Date Accessed: February 10,
2016]
[7] https://en.wikipedia.org/wiki/Activation_energy
APPENDICES
NaOH Standardization
Run
1
2
Mass KHP (g)
0.68
0.67
Vol. NaOH (mL)
11.26
11.23
Concentration of NaOH = [mass KHP(molecular weight
KHP)(1 mol NaOH/1 mol KHP)]/vol NaOH in L
C1=0.298 N
C2=0.294 N
Ave. Concentration = 0.296 N
Reaction Rate
slope ( log c vs t )=
kV
2.303
k =0.022(0.2137)/(2.303)
k =0.00204 /s
1
T
−Ea
log ¿=
R
slope ¿
k vs
Ea =420.66 ( 8.314 51 )
Ea =3497.58
Activation Energy
J
mol