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Experiment No. 2
Rates of Chemical Reactions: Iodination of Acetone
Aguilera, Horacio Miguel N.
11-K
Date Performed: December 10, 2020
Date Submitted: December 21, 2020
ABSTRACT

The goal of this experiment was to determine the rate law of the given reaction mixtures in the iodniation of acetone. The method that was used to obtain the said reactants in their respective experiments was to mix different quantities of 4.0M of acetone, 1.0 M HCl, 0.0050 M iodine, and distilled water. The amount of time in seconds that it took for iodine to lose its color was then documented for each trial. Afterwhich, four experiments with 2 trials each were performed in order to identify the temperature of each reaction mixture, the results per experiment were recorded in the table below
Table 1.0: Rate Laws of the Reaction Mixtures

After the required data gathering was done through the observation of the reactants it was found that the standard deviation of the rate constant was 2.6 × 10^(-6) In scientific notation, alternatively written as 0.0000026. Lastly, the reliability of the gathered data and computations were checked through the standard deviation of k, with the consideration that due to imperfections and therefore a non-ideal environment, along with the possibility for a transcription error, the computed value may not be 100 accurate. Using the previous statement in mind to justify any possible errors, the conducted study resulted in a computed standard deviation of 2.6 x 10-6 or 0.0000026 . This proves the conducted study to be reliable due to the result being relatively accurate since low standard deviation denotes high precision, and therefore a much higher rate of accuracy than that of a higher deviation.
INTRODUCTION
The goal of this experiment was to determine the rate law of the given reaction mixtures in the iodination of acetone. The goal of this experiment was to determine the rate law of the given reaction mixtures in the iodniation of acetone. The method that was used to obtain the said reactants in their respective experiments was to mix different quantities of 4.0M of acetone, 1.0 M HCl, 0.0050 M iodine, and distilled water in a properly sanitized Erlenmayer flask. With the use of a stopwatch, the amount of time in seconds that it took for iodine to lose its color was then documented for each trial.Note that trials were done in thesame room, and therefore have a similar if not thesame environmental condition as a constant variable in this experiment. With a limiation being an uncertainty ;due to a temperature monitoring system not being utilized , a possible room for discrepancy could be implied. The initial concentrations of acetone, I2 , and H+ were then determined for each mixture.Afterwhich, the rate of the reaction was obtained by dividing the initial concentration of I2 by the time needed for I2 to lose color.

From there, the reaction orders of each substance were determined, followed by the rate constant and its standard deviation. Based on these calculations, the rate law for each reaction as well as the reliability of the results in the experiment were identified. The data was gathered and recorded in the form of the table below

Mixture Volume of Acetone in mL 4.0M Volume of HCl in mL 1.0M Volume of I2 in mL 0.0050M Volume of H2O in mL Total Volume of the Solution
1 10.00 mL 10.00 mL 10.00 mL 25.00 mL 55.00 mL
2 5.00 mL 10.00 mL 10.00 mL 30.00 mL 55.00 mL
3 10.00 mL 5.00 mL 10.00 mL 30.00 mL 55.00 mL
4 10.00 mL 10.00 mL 5.00 mL 30.00 mL 55.00 mL
Table 1.1: Substance Volumes

Since the conducted trials were instrumental to obtain the rate law , the necessary equation to compute for it must first be identified. The rate law is a mathematical equation expressing the relationship between the rate of a chemical reaction and the concentration of its reactants (OpenStax, n.d.).



Refer to the equation : aA + bB → products (OpenStax, n.d.)
Equation 1

where a and b = stoichiometric coefficients (OpenStax, n.d.)
A and B= reactants (OpenStax, n.d.)

With a and b being stoichiometric coefficients and A and B being reactants, we may obtain the necessary equation or rate law.(OpenStax, n.d.)

rate = k[A]m[B]n (OpenStax, n.d.)
Equation 2
Referring to equation 2, otherwise known as the Rate Law or Equation,k is the rate constant, while [A] and [B] represent the molar concentrations of the reactants, with the exponents m and n representing the reaction orders of each of the reactants. (OpenStax, n.d.).



A zeroth order reaction has a constant rate, unaffected by the concentration of the reactant (Helmenstine, 2019). A first order reaction has a rate proportional to the concentration of one of the reactants (Helmenstine, 2019). Finally, a second order rate has a rate proportional to the square of the concentration of a reactant (Helmenstine, 2019). The overall reaction order is determined by adding all the orders of each reactant (OpenStax, n.d.).
When performing the experiment, it was also important to know that the rate of the reaction can vary with temperature (“Rates of Chemical,” n.d.). To be more specific, an increase in temperature can increase the reaction rate (“Rates of Chemical,” n.d.). Therefore, it was important to keep a constant temperature across experiments to ensure that results were more accurate. Another important detail to remember during the experiment was that value for the rate constant and the reaction orders had to be determined experimentally by observing the change in reaction rate as the concentration of the reaction changed in the experiment (OpenStax, n.d.).
Knowing this information, the initial concentrations of the reactants were solved in order to complete the rate law. To do this, the formula for initial concentration was used. The formula for initial concentration is simply the molarity formula. Hence, the initial concentration is equal to the moles of solute divided by the liters of solution (Flowers, Theopold, & Langley, et al., 2020):
M = Initial Concentration = (moles of solute)/(liters of solution)
Equation 3
After determining the initial concentration of each reactant, the reaction rate was then determined. Reaction rate, also known as the rate of reaction, refers to the measure of change in concentration of the reactants per unit time (Odufalu, Chacha, Mudda, & Iskandar, 2020). Thus, if the reaction rate for A in Equation 1 was to be determined, it would have this equation (Odufalu, et al., 2020):
rate = – 1/a (∆[A])/∆t
Equation 4
The negative sign is added because the concentration of a reactant always decreases over time (Odufalu, et al., 2020). In the experiment performed, I2 was the substance used to measure the rate of reaction. This is because I2 has a very distinct yellow color that disappears over time as it reacts with acetone, HCl, and water. Thus, the time taken for I2 to react can be easily recorded and used to calculate the reaction rate. Moreover, I2 is always zeroth order in the iodination of acetone (Meyer & Lask, 2010). Hence, the reaction rate of I2 in the experiment was represented by this equation:
rate = – (∆[I_2])/∆t = ([I_2 ]_0)/t
Equation 5
After calculating for this in the experiment, the rate constant and reaction orders were the last remaining unknowns in the equation. The reactions orders were determined by dividing the rate laws by one another. Doing so cancelled out the initial concentrations of certain reactants and allowed the reaction orders to be calculated using logarithms. Once all the reaction orders were identified, the rate constant was then solved for.
Finally, the standard deviation of the rate constant was used to evaluate the reliability of the overall experiment. Since this experiment used the same substances in only slightly varying amounts, the rate constants of each mixture were expected to be close in value to one another. Thus, a lower standard deviation indicated that the rate constant was more consistent between experiments, while a higher standard deviation indicated that the rate constant was less consistent between experiments. In short, the lower the standard deviation of the rate constant, the more reliable the experiment results were. Thus, the sample standard deviation formula was used in the experiment to evaluate its accuracy (Taylor, 2020):
s = √((∑_(i=1)^N▒〖(〖x_i – x ̅)〗^2 〗)/(N –1))
Equation 5
s = sample standard deviation
N = the number of observations
xi = each of the values of the data
(x ) ̅= the mean value of the data set

EXPERIMENTAL
The materials that the experiment made use of were a 125-mL Erlenmeyer flask, three beakers, three watch glasses, a thermometer, a 10-mL pipet, a timer, and a graduated cylinder. a4.0 M acetone, 1.0 M HCl, 0.0050 M iodine, and distilled water. Each of the 3 beakers, prior to being covered with watch glasses were used to contain one of the required amount of the following substances; the acetone, the HCl and iodine respectively, this being done before the experiment proper. To reiterate, the beakers were each covered with a watch glass to prevent entry of any unwanted foreign substances that may potentially lead to an inaccurate result.
After the initial preparations, the first experiment was conducted. It required the use of the a 10-mL pipet in order to transfer 10.00 mL of 4.0 M acetone into a 125-mL Erlenmeyer flask. Thesame 10-mL pipet was then used to transfer 10.00 mL of 1.0 M HCl into the same Erlenmeyer flask that was also containing the acetone. The 25.00 mL of distilled water and the 10.00 mL of 0.0050 M iodine were both measured separately using a graduated cylinder, afterwhich , they were added into thesame Erlenmeyer flask containing all the previously added substances( acetone and HCl). A stopwatch was then promptly set to measure how long it would take until the color of the iodine fully dissapears . Once the the color was completely gone, the timer was stopped. Afterwhich, the temperature of the mixture was measured with a thermometer.
Experiment 2, 3 , and 4 also followed thesame protocol, with the exception being that the quantity of substances was altered between the experiments. The second experiment having 5.00 mL of acetone added, The third having 5.00mL of HCl and fourth having 5.00mL of iodine added into the experiments respectively. Each of the conducted experiments had two trials each, both following thesame protocol. Computations based on the collected data were then calculated for, to reiterate, this was instrumental for the identification of the rate law of the reaction mixtures.

RESULTS AND DISCUSSION
When substituted into the rate law, the equation for this reaction mixture was:
rate = k[C3H6O]m[I2]n[H+]p
To determine the rate law of each reaction in the given experiments, the initial concentration of each reactant was needed. Thus, the moles of solute and liters of solution in each reactant needed to be identified. The table below depicts the moles of solute and liters of solution obtained for each mixture:



Table 1: Moles of Solute and Liters of Solution
Experiment Reactant Moles of Solute (mol) Liters of Solution (L)
1 acetone 0.040 0.055
I2 5.0 × 10-5
H+
0.010
2 acetone 0.020 0.055
I2 5.0 × 10-5
H+
0.010
3 acetone 0.040 0.055
I2 5.0 × 10-5
H+ 5.0 × 10-3
4 acetone 0.040 0.055
I2 2.5 × 10-5
H+ 0.010

These values were obtained using solution stoichiometry. In other words, the moles of solute were obtained by converting the milliliters of solution into liters, then multiplying by the molarity of the substance. Meanwhile, the liters of solution were obtained by adding all of the milliliters of each substance, then converting it into liters, as shown below:
Experiment 1
Moles of acetone = 10.00 mL acetone ((1 L)/(1000 mL)) ((4.0 mol acetone)/(1 L)) = 0.040 mol acetone
Moles of I2 = 10.00 mL I2 ((1 L)/(1000 mL)) ((0.0050 mol I_2)/(1 L)) = 5.0 × 10-5 mol I2
Moles of H+ = 10.00 mL HCl ((1 L)/(1000 mL)) ((1.0 mol HCl)/(1 L))((1 mol H^+)/(1 mol HCl)) = 0.010 mol H+
Liters of Solution = 10.00 mL + 10.00 mL + 10.00 mL + 25.00 mL = 55 mL ((1 L)/(1000 mL)) = 0.055 L
Experiment 2
Moles of acetone = 5.00 mL acetone ((1 L)/(1000 mL)) ((4.0 mol acetone)/(1 L)) = 0.020 mol acetone
Moles of I2 = 10.00 mL I2 ((1 L)/(1000 mL)) ((0.0050 mol I_2)/(1 L)) = 5.0 × 10-5 mol I2
Moles of H+ = 10.00 mL HCl ((1 L)/(1000 mL)) ((1.0 mol HCl)/(1 L))((1 mol H^+)/(1 mol HCl)) = 0.010 mol H+
Liters of Solution = 5.00 mL + 10.00 mL + 10.00 mL + 30.00 mL = 55 mL ((1 L)/(1000 mL)) = 0.055 L
Experiment 3
Moles of acetone = 10.00 mL acetone ((1 L)/(1000 mL)) ((4.0 mol acetone)/(1 L)) = 0.040 mol acetone
Moles of I2 = 10.00 mL I2 ((1 L)/(1000 mL)) ((0.0050 mol I_2)/(1 L)) = 5.0 × 10-5 mol I2
Moles of H+ = 5.00 mL HCl ((1 L)/(1000 mL)) ((1.0 mol HCl)/(1 L))((1 mol H^+)/(1 mol HCl)) = 5.0 × 10-3 mol H+
Liters of Solution = 10.00 mL + 5.00 mL + 10.00 mL + 30.00 mL = 55 mL ((1 L)/(1000 mL)) = 0.055 L
Experiment 4
Moles of acetone = 10.00 mL acetone ((1 L)/(1000 mL)) ((4.0 mol acetone)/(1 L)) = 0.040 mol acetone
Moles of I2 = 5.00 mL I2 ((1 L)/(1000 mL)) ((0.0050 mol I_2)/(1 L)) = 2.5 × 10-5 mol I2
Moles of H+ = 10.00 mL HCl ((1 L)/(1000 mL)) ((1.0 mol HCl)/(1 L))((1 mol H^+)/(1 mol HCl)) = 0.010 mol H+
Liters of Solution = 10.00 mL + 10.00 mL + 5.00 mL + 30.00 mL = 55 mL ((1 L)/(1000 mL)) = 0.055 L
These values were then used to calculate for the initial concentration for the reactants in each experiment:
Initial Concentration of Reactants
Experiment 1
Acetone = (0.040 mol )/(0.055 L) = 0.73 M

I2 = (5.0 × 10^(-5) mol)/(0.055 L) = 9.1 × 10-4 M

H+ = (0.010 mol)/(0.055 L) = 0.18 M
Experiment 2
Acetone = (0.020 mol )/(0.055 L) = 0.36 M

I2 = (5.0 × 10^(-5) mol)/(0.055 L) = 9.1 × 10-4 M

H+ = (0.010 mol)/(0.055 L) = 0.18 M
Experiment 3
Acetone = (0.040 mol )/(0.055 L) = 0.73 M

I2 = (5.0 × 10^(-5) mol)/(0.055 L) = 9.1 × 10-4 M

H+ = (5.0 × 10^(-3) mol)/(0.055 L) = 0.091 M Experiment 4
Acetone = (0.040 mol )/(0.055 L) = 0.73 M

I2 = (2.5 × 10^(-5) mol)/(0.055 L) = 4.5 × 10-4 M

H+ = (0.010 mol)/(0.055 L) = 0.18 M

After calculating the initial concentrations of each reactant, the average rate of reaction was then calculated for each experiment. To do this, the initial concentration of iodine was divided by the average time of each reaction. It is important to take note that milliseconds were converted into seconds by dividing it by 1000.
Rate of Reaction
Experiment 1
Time of Run 1 = 150.05 s
Time of Run 2 = 143.03 s
Average Time = 146.54 s
rate = (9.1 × 10^(-4) M)/(146.54 s) = 6.2 × 10-6 M/s Experiment 2
Time of Run 1 = 282.00 s
Time of Run 2 = 262.00 s
Average Time = 272.00 s
rate = (9.1 × 10^(-4) M)/(272.00 s) = 3.3 × 10-6 M/s

Experiment 3
Time of Run 1 = 308.00 s
Time of Run 2 = 305.02 s
Average Time = 306.51 s
rate = (9.1 × 10^(-4) M)/(306.51 s) = 3.0 × 10-6 M/s
Experiment 4
Time of Run 1 = 70.04 s
Time of Run 2 = 70.07 s
Average Time = 70.06 s
rate = (4.5 × 10^(-4) M)/(70.06 s) = 6.4 × 10-6 M/s




From these reaction rates, a graph was constructed.

Figure 1: Rate of Reaction of Each Experiment
From this graph, it can be seen that the average rate of reaction for second and third experiment were lower than that of the first and fourth experiment. This was attributed to the fact that I2 took a significantly longer time to change color in the second and third experiments when compared first and fourth experiments. Based on this, it was inferred that a slower reaction resulted in a lower rate of reaction.
In addition to this, the information in this graph showed that the average rates of reaction of the first and second experiments were slightly higher than that of the third and fourth experiments:
Average of the Rates of the Reaction
Experiment 1 & 2
Reaction Rate of Experiment 1 = 6.2 × 10-6 M/s

Reaction Rate of Experiment 2 = 3.3 × 10-6 M/s

Average = ((6.2 × 10^(-6) M/s) + (3.3 × 10^(-6) M/s))/2 =
4.8 × 10-6 M/s Experiment 3 & 4
Reaction Rate of Experiment 3 = 3.0 × 10-6 M/s

Reaction Rate of Experiment 4 = 6.4 × 10-6 M/s

Average = ((3.0 × 10^(-6) M/s) + (6.4 × 10^(-6) M/s))/2 =
4.7 × 10-6 M/s

This can most likely be attributed to the slightly lower temperature for third and fourth experiments. In the first and second experiments, the reaction took place in 27"℃" while third and fourth experiments took place in 26.5"℃" . It was deduced that since the temperature was lower in third and fourth experiments, the rate was lowered as well. Based on this, it was concluded that the results of the experiments may not be completely accurate due to the temperature changes that occurred in between experiments.
However, despite the potential discrepancies that the data might have, the computed reaction rates were still substituted into the rate law along with the calculated initial concentrations. These results were yielded:

Experiment 1
6.2 × 10-6 M/s = k[0.73 M]m[9.1 × 10-4 M]n[0.18 M]p
Experiment 2
3.3 × 10-6 M/s = k[0.36 M]m[9.1 × 10-4 M]n[0.18 M]p
Experiment 3
3.0 × 10-6 M/s = k[0.73 M]m[9.1 × 10-4 M]n[0.091 M]p
Experiment 4
6.4 × 10-6 M/s = k[0.73 M]m[4.5 × 10-4 M]n[0.18 M]p

This left only k, m, n, and p as the unknowns in the rate law. Thus, the reaction orders (m, n, and p) were solved by dividing the reaction rates by one another.
Reaction Orders
Reaction Order of Acetone (m)
Using experiment 1 and 2:

(6.2 ×10^(-6) M/s)/(3.3 ×10^(-6) M/s) = (k[0.73 M]_1^m [9.1 ×10^(-4) M]_1^n [0.18 M]_1^p)/(k[0.36 M]_2^m [9.1 ×10^(-4) M]_2^n [0.18 M]_2^p )

(6.2 ×10^(-6) M/s)/(3.3 ×10^(-6) M/s) = (k[0.73 M]_1^m [9.1 ×10^(-4) M]_1^n [0.18 M]_1^p)/(k[0.36 M]_2^m [9.1 ×10^(-4) M]_2^n [0.18 M]_2^p )
1.9 = (2.0)m
log(1.9) = mlog(2.0)
m = (log⁡(1.9))/(log⁡(2.0))
m = 0.93 = 1 Reaction Order of I2 (n)
Using experiment 1 and 4:

(6.4 ×10^(-6) M/s)/(6.2 ×10^(-6) M/s) = (k[0.73 M]_4^m [4.5 ×10^(-4) M]_4^n [0.18 M]_4^p)/(k[0.73 M]_1^m [9.1 ×10^(-4) M]_1^n [0.18 M]_1^p )

(6.4 ×10^(-6) M/s)/(6.2 ×10^(-6) M/s) = (k[0.73 M]_4^m [4.5 ×10^(-4) M]_4^n [0.18 M]_4^p)/(k[0.73 M]_1^m [9.1 ×10^(-4) M]_1^n [0.18 M]_1^p )
1.0 = (0.49)n
log(1.0) = nlog(0.49)
n = (log⁡(1.0))/(log⁡(0.49))
n = 0
Reaction Order of H+ (p)
Using experiment 1 and 3:

(6.2 ×10^(-6) M/s)/(3.0 ×10^(-6) M/s) = (k[0.73 M]_1^m [9.1 ×10^(-4) M]_1^n [0.18 M]_1^p)/(k[0.73 M]_3^m [9.1 ×10^(-4) M]_3^n [0.091 M]_3^p )
(6.2 ×10^(-6) M/s)/(3.0 ×10^(-6) M/s) = (k[0.73 M]_1^m [9.1 ×10^(-4) M]_1^n [0.18 M]_1^p)/(k[0.73 M]_3^m [9.1 ×10^(-4) M]_3^n [0.091 M]_3^p )
2.1 = (2.0)p
log(2.1) = plog(2.0)
p = (log⁡(2.1))/(log⁡(2.0))
p = 1.1 = 1

Based on these calculations, a table was constructed:
Table 2: Reaction Orders of Each Reactant
Reactant Reaction Order
acetone 1
I2 0
H+ 1

From this table, it can be said that acetone and H+ were first order while I2 was zeroth order. Therefore, the overall reaction order was second order. It can also be said that the concentration of acetone and H+ was directly proportional to the rate of the reaction. This can be observed when comparing the reaction rates and initial concentrations in the first and second experiments, as well as the first and third experiments.
Experiment 1 and Experiment 2
Reaction Rate of Experiment 1 = 6.2 × 10-6 M/s
Reaction Rate of Experiment 2 = 3.3 × 10-6 M/s Conc. of Acetone in Experiment 1 = 0.73 M
Conc. of Acetone in Experiment 2 = 0.36 M

Experiment 1 and Experiment 3
Reaction Rate of Experiment 1 = 6.2 × 10-6 M/s
Reaction Rate of Experiment 3 = 3.0 × 10-6 M/s Conc. of H+ in Experiment 1 = 0.18 M
Conc. of H+ in Experiment 3 = 0.091 M

When viewing these comparisons, it was noticed that reaction rate doubled when the concentration of acetone doubled. Similarly, the rate of reaction doubled when the concentration of H+ was doubled. This supports the idea that acetone and H+ were first order reactions.
As for I2, it was observed in the experiment that the concentration of I2 had relatively no effect on the rate of reaction. This can be seen when comparing experiments 1 and 4.
Experiment 1 and Experiment 4
Reaction Rate of Experiment 1 = 6.2 × 10-6 M/s
Reaction Rate of Experiment 4 = 6.4 × 10-6 M/s Conc. of I2 in Experiment 1 = 9.1 × 10-4 M
Conc. of I2 in Experiment 4 = 4.5 × 10-4 M

The rate of the reaction barely changed when the concentration of I2 was changed, indicating that I2 was a zeroth order reaction. Thus, it can be said that the data in Table 2 was proven experimentally.
Finally, since the reaction order had been determined, it was concluded that the rate laws for each mixture were:
Rate Laws of Each Experiment
Experiment 1
6.2 × 10-6 M/s = k[0.73 M]1[9.1 × 10-4 M]0[0.18 M]1
Experiment 2
3.3 × 10-6 M/s = k[0.36 M]1[9.1 × 10-4 M]0[0.18 M]1
Experiment 3
3.0 × 10-6 M/s = k[0.73 M]1[9.1 × 10-4 M]0[0.091 M]1
Experiment 4
6.4 × 10-6 M/s = k[0.73 M]1[4.5 × 10-4 M]0[0.18 M]1
From these rate laws, the rate constant of each reaction mixture was then determined.

Rate Constant of Each Experiment
Experiment 1
6.2 × 10-6 M/s = k[0.73 M]1[9.1 × 10-4 M]0[0.18 M]1
k = (6.2 × 10^(-6) M/s )/([0.73 M]^1 [9.1 ×10^(-4) M]^0 [0.18 M]^1 )
k = 4.7 × 10-5 1/Ms
Experiment 2
3.3 × 10-6 M/s = k[0.36 M]1[9.1 × 10-4 M]0[0.18 M]1
k = (3.3 × 10^(-6) M/s )/([0.36 M]^1 [9.1 ×10^(-4) M]^0 [0.18 M]^1 )
k = 5.1 × 10-5 1/Ms
Experiment 3
3.0 × 10-6 M/s = k[0.73 M]1[9.1 × 10-4 M]0[0.091 M]1
k = (3.0 × 10^(-6) M/s )/([0.73 M]^1 [9.1 ×10^(-4) M]^0 [0.091 M]^1 )
k = 4.5 × 10-5 1/Ms
Experiment 4
6.4 × 10-6 M/s = k[0.73 M]1[4.5 × 10-4 M]0[0.18 M]1
k = (6.4 × 10^(-6) M/s )/([0.73 M]^1 [4.5 ×10^(-4) M]^0 [0.18 M]^1 )
k = 4.9 × 10-5 1/Ms

Obtaining the average of these values yielded the average value of the rate constants in the reaction mixture:
Average Value of the Rate Constants
Average Value of k = ((4.7 × 10^(-5) 1/Ms) + (5.1 × 10^(-5) 1/Ms) +(4.5 × 10^(-5) 1/Ms) + (4.9 × 10^(-5) 1/Ms))/4
Average Value of k = 4.8 × 10-5 1/Ms
From this data, a graph was constructed:

Figure 2: Values of the Rate Constant
This graph compares the value of the rate constant from each experiment to one another and to their average. In this graph, it can be observed that the rate constant for each experiment were relatively close in value to one another and to the average. This indicated that the rate constants were reasonably consistent in value with one another. Thus, to further analyze the consistency between the rate constants, the standard deviation of the rate constant was obtained using the sample standard deviation formula.
s = √((∑_(i=1)^N▒〖(〖x_i – x ̅)〗^2 〗)/(N –1))
s = √((2.0 × 10^(-11))/(4 –1))
s = 2.6 × 10^(-6)
Therefore, the standard deviation of the rate constant was 2.6 × 10^(-6). This number is very close to zero. Therefore, it indicated that the values of the rate constants in the experiment were close to the average value of the rate constant, which meant that their values were also fairly consistent with one another. Based on these results, it was concluded that the experiment was mostly accurate and precise, despite the temperature changes between experiments.

CONCLUSION
The initial concentration of each reaction misture was identified using the volume and molarities, afterwhich, with iodine used as a dependent variable in each experiment, the reaction rate of each experiment was computed for . The results of the experiment show that possibly due to the fact the time that iodine took to change color in the second and third experiments than the latter , and since a slow reaction time denotes a lower reaction rate. the reaction rate of the second and third experiments were of a lower quantity than the first and fourth. Further observation showed that due to the average


In addition to this, it was found that the average of the reaction rates for the first and second experiments were slightly higher than that of third and fourth experiments. This was attributed to the slight decrease in temperature for the third or fourth experiment, which could have resulted in a lower reaction rate. Hence, the experiments might not have been completely accurate considering that the change in temperature could affected the results of the data. Thus, experiments involving rates of reaction should always make an effort to keep a constant temperature throughout the entire experimentation process to avoid potentially inaccurate results.
Afterwards, the reaction orders for each substance were solved. It was found that acetone and H+ were first order reactions while I2 was a zeroth order reaction. Hence, it was concluded that the reaction rate increased proportionately to the concentration of acetone and H+. Meanwhile, the reaction rate was unchanged when I2 increased or decreased. Moreover, it was concluded that the overall reaction order was second order.
Based on the information gathered, the rate laws of the experiment were then able to be identified:
Rate Laws of Each Experiment
Experiment 1
6.2 × 10-6 M/s = k[0.73 M]1[9.1 × 10-4 M]0[0.18 M]1
Experiment 2
3.3 × 10-6 M/s = k[0.36 M]1[9.1 × 10-4 M]0[0.18 M]1
Experiment 3
3.0 × 10-6 M/s = k[0.73 M]1[9.1 × 10-4 M]0[0.091 M]1
Experiment 4
6.4 × 10-6 M/s = k[0.73 M]1[4.5 × 10-4 M]0[0.18 M]1
After identifying the rate laws, the rate constant of each experiment was identified. It was found that the average of the rate constant was 4.8 × 10-5 1/Ms. Furthermore, it was also found that the standard deviation between the rate constants was 2.6 × 10^(-6). Based on this, it was determined that the values of the rate constants in the experiment were close to the average rate constant, indicating that their values were also reasonably consistent with one another. Therefore, even though the temperatures between experiments varied slightly, the experiments were still concluded to be mostly accurate and precise.












































REFERENCES

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