CHM 431 PHYSICAL CHEMISTRY
TITLE
: KINETICS- ORDER OF REACTION
NAME
: NURFADHILAH BINTI JAAFAR
ID NUMBER
: 2016675256 2016675256
GROUP
: AS 246 3B
LAB PARTNER
: SITI FATIMAH BINTI SHAIKH ABDUL MUNAIM
SUBMITION DATE
:8
LECTURER’S NAME
: MADAM ZARILA MOHD SHARIFF
TH
NOVEMBER 2016
TITLE Kinetics- Order of Reaction OBJECTIVES a) To determine the order of a reaction INTRODUCTION The study of rates of chemical process is related to chemical kinetics which is a branch of chemistry and also known as reaction kinetics. The rate of homogenous reaction at any instant is proportional to the product of the molar concentration of the reactants raised to power determined from the experiment. The order of the reaction can be express on the following form: Rate= k[A] x [B]y[C]z The letters A, B and C represent the concentration of the reactants. k is referring the rate constant, while the exponents x, y and z are called the rate law of exponents or also known as the orders of the reaction. The methods of initial rates were used, in which the reaction of interest is carried out at different concentrations and the rate of trial is measured. They do not change in temperature and must be determined experimentally and the temperature must hold at constant during the kinetic experiment. The magnitude of the rate constant indicates the speed of the reaction. Hence, a small rate constant indicates a slow reaction, while large rate constant indicates a faster reaction. The experiments that need to be observed are between potassium permanganate and oxalic acid. A deep purple colour of potassium permanganate will change to light brown when react to oxalic acid and it can be observed visually. The rate of reaction can be described by the general equation called the rate law equation: Rate= k[KMnO 4]x [H2C2O4]y The Rate Method of Initial Rates The initial rates method is a way to make the reaction of a reactant effectively constant. Initial rate of a reaction is defined as the rate of reaction at the instant the reactants are first mixed. The time it takes to use up the permanganate and the speed are the subjects that need to be focused to determine the speed of the reaction.
Rate =-d[KMnO4]/dt= {[KMnO 4]final –[ KMnO4]initial} (tfinal – tinitial) Since [KMnO 4]final = 0, and if we set t initial = 0, the above equation simplifies to: Rate= -(-[KMnO4]initial/tfinal) = [KMnO 4] /t The exponents x and y can find by using the method of initial rates. It is impossible to make the measurements from the exact instant a reaction starts because the initial rates are always estimation. Hence the method involves the measurement and comparison of the initial rates of reaction when different initial concentrations are used. The initial rate for each being potassium permanganate concentration divided by time. The rate of expressions for experiment 1 and experiment 2 can be written as follow: Rate1 = k[ KMnO 4]1 x[H2C2O4]1y
Rate1 = [KMnO4]1/t1
Rate2 = k [KMnO4] 2 x[H2C2O4]2y
Rate2 = [KMnO4]2/t2
Taking the ratio of the two rates: Rate1/Rate2 =([ KMnO 4]1/[KMnO 4]2)x (Note that in these two reactions [H 2C2O4]1 = [H2C2O4]2, thus can be cancelled out). Thus we can determine x. In a similar manner, y, the order of reaction with respect to oxalic acid can be determined if we compare experiment 1 and 3)
PROCEDURE 1. Three burettes were set up; one contained 0.02 M KMnO 4, one with 0.5M H 2C2O4 and one with distilled water. A label was placed on each because it was difficult to differentiate between water and oxalic acid. The exact molarities of the KMnO 4 and H2C2O4. 2. The required amount of H 2C2O4 and distilled water was place into a thoroughly washed and dried conical flask. The amounts are dictated by the experiment that you are doing as Table 4.1. When we was overshoot we need to start again. This lab is very dependent on dispensing the exact quantities.
Reagents
Exp 1 (cm 3)
Exp 2 (cm 3)
Exp 3 (cm 3)
H2C2O4
20
20
10
KMnO4
10
5
10
H2O
0
5
10
Table 4.1 Volume of reagents required 3. The required amount of KMnO 4 was placed into a 15 cm test tube. 4. Permanganate was added to the oxalic acid and the timing was started when the permanganate tube was emptied. The conical flask was mixed thoroughly by swirling until the solution turns a light yellow/ brown colour. The timing was stopped and the time it actually took for the reaction to take place was recorded. 5. The experiment was repeated with a second and third trial. The average of these 3 was took as the reaction time. 6. Steps 3 through 5 were repeated for experiments 2 and 3. 7. The rate for each of the three experiments was determined. That was just the [KMnO4]/taverage 8. The full rate equations for each of the 3 experiments were determined. Note that
Rate = k[KMnO 4]x[H2C2O4]y The values of k, x and y were determined.
CHEMICALS 1. Potassium Permangante (KMnO 4), 0.020 M 2. Oxalic acid, (H 2C2O4), 0.50 M
DISCUSSIONS
The experiment was done to observe the order of the reaction of oxalic acid and potassium permanganate by changing their initial molar concentration. The molarity for experiment 1, experiment 2 and experiment 3 for H 2C2O4 and KMnO 4 are 0.50 M and 0.020 M respectively. The volume used for H 2C2O4 and KMnO 4 are 20 ml and 10 ml respectively in experiment 1. For experiment 2, the volume of H 2C2O4 used remain the same, 20 ml, but volume KMnO4 is reduce to 10 ml. Volume for KMnO4 and H 2C2O4 used in experiment 3 both reduce at the same volume which are 10 ml . Each experiment conducted with 3 trials. In experiment 1 the average time it took to completely react are 191s while in experiment 2 it took 218s on average. Experiment 3 took the longer time on average for the solution to change is colour which is 348s. The correct mechanism for this reaction at experiment 3 could be slow since the rate of reaction dependent on the concentration of KMnO 4.
The sample of calculation is shown below:
Calculations for The Concentration of KMnO 4 Experiment 1 and 2 = no dilution the concentration remain 0.50 m Experiment 3 =
0.50 M(0.01) = C f x (0.02) Cf = 0.25 M
Calculations for The Concentration of H 2C2O4 Experiment 2: CiVi = Cf Vf (0.02 M)(0.005 M) = C f (0.01) Cf = 0.010 M The average time 350 s + 343s + 353 s = 348s 3 The Initial Rate Experiment 1 = Rate 1= [KMnO 4] =0.20 M T1
=1.07 x 10 -4 Ms-1
191s
Rate 1 = [H 2C2O4] = 0.50 M =2.68 x 10 -3 Ms-1 T1
186.6 s
The order of the reaction was determined by one of the common method which is method of initial rates. The initial concentration for each solution is being consecutively changed to observe how these factors can affect the rate of reaction. In experiment 1 and 2 the initial concentration of KMnO 4 is reduced 2 times cause the reaction rate to decrease by factor 4.93 x10 -5. As shown in calculation the reaction order that we get is 2, but supposedly we should get the reaction of order is x =1. The reaction should be First Order with respect to KMnO 4. Now we compare Expts 1 and 3. Here a doubling of the initial KMnO4 concentration causes no change in the Reaction Rate. Hence, making the overall reaction order y = 2 and the reaction is Second Order with respect to H 2C2O4 supposedly. Thus the rate of law of reaction should be: Rate: k[KMnO 4]1[H2C2O4]2 The average rate of constant was calculated to be 128.66 s -1. There were several potential errors that happen in conducting the experiment that cause for us to not getting the expected result of the order of the reaction. Some errors could include failing to dispense the exact amount of reactant into the test tube, failing to mix the solution after each reactant or failing to time properly for the each of the reaction. For instance, when the timer was not stopped when the solution reached the desired colour change, it could alter the results, because each reaction was carried out to different point, not necessarily the end point. All these errors could affect the rate of reaction from getting the expected results.
CONCLUSIONS
In conclusion, we could not get the expected result of the order of the reaction due to some errors. The expected order of the reaction should be Rate: k[KMnO 4]1[H2C2O4]2 . The average times taken for the reaction to takes place for experiment 1, experiment 2 and experiment 3 are 191s, 218 s and 348s. On the other hand, the rate reaction for experiment 1, experiment 2 and experiment 3 are 2.68 x 10 -3 and 4.93 x 10 -5 and 7.59 x 10 -4 Ms-1. However the average k is 128.66 s -1
QUESTIONS
1. In this experiment the orders of reaction x and y are obtained by taking ratios of rates for two trials or experiments. a) Explain how this experiment can be modified so that the orders of reaction can be obtained by plotting appropriate graphs. Graphical method can be also be applied to second order of the reactions. A plot of 1/[A]t versus t gives a straight line with a slope of k.
b) What plots must be done and explain how the orders of reaction can be determined from the plots. Use the data in the table to separately plot concentration, the natural logarithm of the concentration, and the reciprocal of the concentration (the vertical axis) versus time (the horizontal axis). Compare the graphs with those in Figure 14.16 "Properties of Reactions That Obey Zeroth-, First-, and Second-Order Rate Laws" to determine the reaction order .Write the rate law for the reaction. Using the appropriate data from the table and the linear graph corresponding to the rate law for the reaction, calculate the slope of the plotted line to obtain the rate constant for the reaction. For zero-order reactions, graph concentration vs. time to get a line with the slope -k. For first-order reactions, graph the logarithm of concentration vs. time to get a line with the slope -k. For second-order reactions, graph the reciprocal of concentration vs. time to get a line of slope -k.
c) Would this method (graphically) be more accurate than what has been done experimentally? Yes.
It is because
plotting the concentration of a reactant as a function of time
produces a graph with a characteristic shape that can be used to identify the reaction order in that reactant.
2. Orders of reaction are normally integers. Is it possible to have non-integers. E.g fractions as order of the reaction? if yes give an example of such reaction. Yes there are reaction inn fractional order reactions, the order is a non-integer, which often indicates a chemical chain reaction or other complex reaction mechanism For example, the pyrolysis of ethanal (CH3CHO) into methane and carbon monoxide
proceeds with an order of 1.5 with respect to ethanal: r = k[CH3CHO]3/2.[13] The decomposition of phosgene (COCl2) to carbon monoxide and chlorine has order 1 with respect to phosgene itself and order 0.5 with respect to chlorine: r = k[COCl2] [Cl2]1/2.
REFERENCES
1. Iskandar, A. (2012, April 21). Reaction Order. Retrieved October 25, 2016, from http://chem.libretexts.org/Core/Physical_and_Theoretical_Chemistry/Kinetics/Rate_Laws /The_Rate_Law/Reaction_Order 2. Chang, R. (2005). Physical Chemistry for the Biosciences. Sansalito, CA: University Science 3. Launer, H. F., & Yost, D. M. (1934). The Kinetics of the Reaction between Potassium Permanganate and Oxalic Acid. II. Journal of the American Chemical Society, 56(12), 2571-2577. doi:10.1021/ja01327a013
