Kintetics-Chemical Reaction Rates Practical PDF

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General Chemistry II Kinetics Lab...

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1 Roseanne Power Chamberlin Dr. Tina Ross General Chemistry II Lab Date of Experiment - 6 February 2019

Experiment #3- Chemical Reaction Rates Introduction Kinetics in chemistry is the study of reaction rates. The rate of a chemical reaction depends on several factors such as, 1) the nature of the reaction, 2) the concentration of the reactants, the temperature, and the presence of a possible catalyst (Ebbing and Gammon, 442 443). The objective of this experiment was to study the nature of a reaction rate with respect to concentration, surface area and catalyst. Part I looked at the kinetics of iodination of acetone by changing reactant concentrations in the presence of an aqueous acid that acts as a catalyst and ultimately determining the rate law. Part II studied the effects of (A) Surface area (hydrochloric acid, HCl and marble chips, CaCO3), and (B) Catalyst (manganese oxide and 3% hydrogen peroxide). Theory Part I When an aqueous iodine solution reacts with acetone in the presence of an acid, the yellow color slowly fades as the iodine is consumed. The products of the reaction are iodoacetone and hydrogen iodide. The hydrogen ion is a catalyst for this reaction. The overall

2 stoichiometric of this reaction is (CH3)2C=O + I2 + H+ → (CH2I)C=O + HI. The rate law for this reaction is of the form: Rate=k[acetone]m[H+]n[I2]p presented in (K.J.Laidler, 2004). Here, Iodine is made the “limiting reactant” in a large excess of acetone and H+ ion. The rate of the reaction is determined by measuring the time required for the initial concentration of Iodine I2 to be used up completely by change of color (V.D. Athawle, and Parul Mathur, 2001). Rate=[I2]/time It is established that the rate information is used to determine the order of reaction with respect to acetone (m), acid (n), and iodine(p) by varying the amount of reactants and measuring the effects of rate. Subsequently, the rate constant (k) is obtained once the order of reaction is known. Part II Surface area of a reactant determines how fast or how slow a reaction undergoes. Increased 

surface area increases reaction rate while decreased surface area slows reaction rate. Also, addition of a catalysts speeds up a reaction rate by lowering its activation energy. Methodology Equipment and Materials: The tools and reagents used in this experiment included the following: (1) four 125mL Erlenmeyer flasks; (2) 10 mL graduated cylinder; (3) Four small beakers; (4)Sheet of white paper; (5) Test tube rack, test tubes ( 15 mL); (6)Manganese dioxide; (7) 3-4% hydrogen peroxide; (8) 1 mL Hydrogen chloride ( HCl); (9) 0.005 M iodine; (10) 4 M acetone:

3 Part I: Finding the rate, orders of reaction and rate constant, k at room temperature: Experimental set-up Table 1.0 volume (mL) of Reactants for each iodination Experiment Experiment

HCl (mL)

Acetone(mL)

I2(mL)

Water (mL)

Total (mL)

1

5

5

5

10

25

2

5

5

10

5

25

3

5

10

5

5

25

4

10

5

5

5

25

Four 125mL flasks were initially collected from the lab shelf and placed on the bench to begin the experiment. The experiment was set-up using Table 1.0. The respective molar volumes of each reactants (1.0 M HCl, 4.0 M and 0.0050 iodine) were collected in the 125ml. Experiment one was started with a stopwatch as soon as the iodine solution was added to the solution of acetone and HCl. The solution flask was swirled for as long as the deep yellow color remains. It was swirled against a white paper background to get a clear view of when the yellow color faded. The stopwatch was stopped once the yellow color changed. Time was noted in the data-sheet for kinetics section. Experiment two continued in like manner until all the experiment was completed. Data is presented in Table 2.0. Part II: Surface Area(A) and Catalyst(B) Procedure (A) Sufficient amount of marble CaCO3, chips and ground was added to two separate test  tubes. Next, 2 mL of 6.0 M HCl was added to each tube in which a contrasting observation between the two test tubes was noted in the data-sheet.

4 (B) 5 mL of 3% hydrogen peroxide was added to 3 test tubes labeled (1)boiling chips, (2) catalyst or Manganese (IV) dioxide (MnO2), and (3)control ( without catalyst). Rate of O2 gas generated was noted in data-sheet ( Table 3.0) Data Collection and Analysis Part I: Kinetics Table 2.0 presents rates of disappearance for I2 Experiment

Initial Conc. Acetone (M)

Initial Conc. HCl(M)

Initial Conc. I2(M)

Time (s)

Rate M s-1

1

0.8000

0.2000

0.0010

154.7

6.5x10-6

2

0.8000

0.2000

0.0020

303.12

6.6x10-6

3

1.6000

0.2000

0.0010

120.5

8.3x10-6

4

0.8000

0.4000

0.0010

70.6

1.4x10-5

The rate of disappearance for I2 was determined by using M1V1=M2V2 and divide that by the   time noted for the yellow color to fade. It was interesting to note here for experiment 1) and 2), that when I2 concentration was doubled, it doubled the time of disappearance but not so much the  rate. It had very little change and therefore the order to reaction with respect to I2 is Zero. But the order of reaction with respect to HCl and Acetone is First. See calculations below. Table 3.0 presents Iodination reaction constant, k for each Experiment Experiment

K M-1s-1

1

4.1x10-5

2

4.1x10-5

3

2.6x10-5

4

4.4x10-5

5

Average

3.8x10-5

Standard Deviation(SD)

7.0x10-6

It was interesting to note that k values are fairly constant for all experiments with varying reactant concentrations as expected, because rate constant (k) is independent of concentration. Calculations 1. Final concentration of a)acetone, b)HCl and c)Iodine in experiment 1-4: M1V1=M2V2 Experiment (1) a) Acetone M1V1= M2V2 = (4.0M)(5.00mL)(25mL)(M2)=0.8000M b) HCl M1V1= M2V2= (1.0M)(5.00mL)(25mL)(M2)=0.2000M c) I2 M1V1= M2V2= (0.0050M)(5.00mL)(25mL)(M2)=0.0010M Experiment (2) a) Acetone M1V1= M2V2 = (4.0M)(5.00mL)(25mL)(M2)=0.8000M b) HCl M1V1= M2V2= (1.0M)(5.00mL)(25mL)(M2)=0.2000M c) I2 M1V1= M2V2= (0.0050M)(10mL)(25mL)(M2)=0.0020M Experiment (3) a) Acetone M1V1= M2V2 =  (4.0M)(10mL)(25mL)(M2)=1.6000M b) HCl M1V1= M2V2= (1.0M)(5.00mL)(25mL)(M2)=0.2000M c) I2 M1V1= M2V2= (0.0050M)(5.00mL)(25mL)(M2)=0.0010M Experiment (4) a) Acetone M1V1= M2V2 = (4.0M)(5.00mL)(25mL)(M2)=0.8000M b) HCl M1V1= M2V2= (1.0M)(10mL)(25mL)(M2)=0.4000M c) I2 M1V1= M2V2= (0.0050M)(5.00mL)(25mL)(M2)=0.0010M 2. Rate of Reaction for experiment 1-4 (1) Rate = I2/t = 0.0010M/154.7s = 6.5x10-6  Ms-1 (2) Rate = I2/t = 0.0020M/300.12s = 6.6x10-6  Ms-1

6 (3) Rate = I2/t = 0.0010M/120.5s = 8.3x10-6  Ms-1 (4) Rate = I2/t = 0.0010M/70.6s = 1.4x10-5  Ms-1 3. Reaction Order Rate2 Rate1

p= Rate2 Rate1

p= Rate2 Rate1

p=

=

k[acetone]m [HCl]n [I2]p m n p k[acetone] [HCl] [I2]

log 1.01538 log 2

=

=

=

(0.0020)p p (0.0010)

= 2p 1.01538= 2p

=

8.3x10−6 4.0x10−6

=

(1.600)m m (0.800)

= 2m 2.08= 2m

= 1.053 = 1. T heref ore the order with respect to Acetone is F irst Order.

k[acetone]m [HCl]n [I2]p m n p k[acetone] [HCl] [I2]

log 2.08 log 2

6.6x10−6 6.5x10−6

= 0.02203 = 0. T heref ore the order with respect to I odine is zero.

k[acetone]m [HCl]n [I2]p m n p k[acetone] [HCl] [I2]

log 2.08 log 2

=

=

8.3x10−6 4.0x10−6

=

(0.400)m m (0.200)

= 2m 2.08= 2m

= 1.053 = 1. T heref ore the order with respect to H Cl is F irst Order.

4. Calculate k for experiment 1 to 4 R ate = k [acetone]m [HCl]n [I2]p 1) 6.5x10−6 = k[0.8]1 [0.2]1 [0.0010]0 =

6.5x10−6 [0.8]1 [0.2]1 [0.0010]0

2) 6.6x10−6 = k[0.8]1 [0.2]1 [0.0020]0 =

= k = 4.1x10−5 M −1 s−1

6.6x10−6 [0.8]1 [0.2]1 [0.0010]0

= k = 4 .1x10−5 M −1 s−1

3) 8.3x10−6 = k[1.6]1 [0.2]1 [0.0010]0 =

8.3x10−6 [1.6]1 [0.2]1 [0.0010]0

= k = 2 .6x10−5 M −1 s−1

4) 1.4x10−5 = k[0.8]1 [0.4]1 [0.0010]0 =

1.4x10−5 [0.8]1 [0.4]1 [0.0010]0

= k = 4 .4x10−5 M −1 s−1

5) Average

4.1x10−5 +4.1x10−5 +2.6x10−5 +4.4x10−5 4

= 3.8x10 −5M −1s −1

6) STD (4.1x10−5 − 3.8x10−5 )2 + (4.1x10 −5 − 3.8x10 −5) 2 + ( 2.6x10−5 − 3.8x10−5 )2 + (4.4x10−5 − 3.8x10 −5) 2 = 7) Parts per 1000

SD Av

x1000 =

7.0x10−6 3.8x10−5

1.98x10−10 4

= √ 4.95x10 −11 = 7.0x10 −6

x1000 = 1.84x102 or 184.2 parts per 1000 value f or k

Part II: Surface Area ( A) and Catalyst (B)

7 Table 4.0 presents the results of reaction rates for Part II A and B experiments. Surface Area of marble, CaCO3(s)

Bubbling Rates ( Fast or Slow) Observation

Chips

Slow release of bubbles ( slow CO2 (g) formation)

Ground

Fast ( immediate release of bubbles or formation of CO2): As soon as a HCl drop touches the reactant surface.

Catalyst Chip

Slow release of bubbles - slow O2 (g) formation

Manganese Oxide

Fast - bubbles of O2 (g) release super fast.

Blank Control

No change

The rate of reaction observed for HCl and marble chips was quite faster compared to ground marble and HCl as indicated by the formation of CO2 gas or bubble release. Similar observation was noted for the reaction involving a catalyst. Discussion Part I Iodination reaction of acetone was robust. Experiment four had the fastest reaction rate compared to experiments 1-3. It took only 1.1 minutes or 70 seconds for the yellow color which can be explained by the access amount of concentration of HCl (H+ ) and acetone concentration. Increase concentration of the acid affects the reaction rate. It acts as an catalyst by speeding up the reaction. The effects of varying the concentrations of acetone, iodine and hydrogen ions have been studied earlier ( Siddharth Kulkarni and S. T. Shukla, 2015) and it has been found that the reaction is zero order with respect to iodine which was proven in this experiment. In addition

8 when reaction orders are first with respect to acetone and hcl. This was confirmed yet by another study (Prem D. Sattsangi, 2011). The reaction rate constant was independent as noted for the Iodine concentration which confirms that rate constant in a zero order reaction is not dependent on concentration. Part II Table 4.0 presented the results for the reactions involving surface area and a catalyst. As briefly mentioned above that increased surface area has increases reaction rate which was observed for the reactions between HCL and CaCO3 marble ground compared to reaction of marble chips and HCl. The reactions involving a catalyst was super robust as noted with the immediate release of gas bubbles indicating formation of CO2 gas. Conclusion The experiment was successfully met its objectives in both parts (I and II). It can be concluded that varying the concentration or acetone and hydrogen and limiting iodine concentration leads to a zero order reaction. Although, the overall order of reaction with respect to acetone and HCl is second order. Increased surface area and an addition of a catalyst proved to increase reaction rates. Works Cited 1. Athawle, Parul Mathur ; Experimental Physical Chemistry, New Age Intl. Pvt. Ltd.(2001)79-80. Accessed 23rd Feb. 2019. 2. Ebbing,D.& Gammon,S.D.General Chemistry Text Book ,11th edition CENAGEbrain publication, (2017)442-3.

9 3. Findlay’s practical physical chemistry 9th edition revised and edited by B.P. Levitt Longman Group Ltd (1973) 338-340. Accessed 23rd Feb. 2019. 4. K.J.Laidler,Chemical Kinetics 3rd edition Pearson Education (2004) 392-394. Accessed 23rd Feb. 2019. 5. Prem D. Sattsangi J. Chem. Educ., 2011, 88 (2), pp 184–188. Accessed 23rd Feb. 2019. 6. Shriniwas L Kelkar and Dilip D Dhavale Resonance October (2000) 24-31 Pre-Lab Questions 1. What is a heterogeneous reaction? A chemical reaction in which the reaction are in two or more phases. A reaction that takes place on the surface of a catalyst of a different phase are also called heterogeneous. 2. Why would we expect temperature to affect the rate of a chemical reaction? This can be explained by Arrhenius equation which shows that rate of a reaction is dependent on temperature. And the underlying factor is the frequency factor (A) which relates to frequency of collisions with proper orientation (pz). The frequency factor has a dependency on temperature as described by collision theory. The collision theory assumes that for a reaction to occur, reactant molecules must collide with an energy greater than some minimum value and proper orientation. 3. In 1913 Marie Curie and Heike Kammerlingh Onnes discovered that radium undergoes nuclear decay at the same rate as liquid hydrogen temperature (-2530 C) as at room temperature? What does this suggest about the nature of nuclear decay? T  his suggests that the nature of nuclear decay is dependent on an ideal environment or optimum

10 temperature and it happens very slowly and at a very low temperature as that of liquid hydrogen. Post-Lab Questions 1. a) Molarity of acetone=

M 1V 1 V2

= M2 =

4.0M x5.00mL = 25.mL

0.80 M

b) Rate = I2/t = 0.0010M/250 s = 4.0x10-6  Ms-1 2. a) Molarity of acetone=

M 1V 1 = V2

M2 =

4.0M x10mL = 25.mL

1.600 M

b) Rate = I2/t = 0.0010M/120 s = 8.3 x10-6  Ms-1 c) Order of reaction (m) with respect to acetone Rate2 Rate1

p=

=

k[acetone]m [HCl]n [I2]p m n p k[acetone] [HCl] [I2]

log 2.08 log 2

=

8.3x10−6 4.0x10−6

=

(1.600)m m (0.800)

= 2m 2.08= 2m

= 1.053 = 1. T heref ore the order with respect to Acetone is F irst Order.

3. Molarity of I2=

M 1V 1 V2

= M2 =

0.005M x10mL 25.mL

= 0.0020 M

Rate = I2/t = 0.0020M/ = 8.3 x10-6  Ms-1, t t = 0.0020M ÷ 8.3 x10−6 M −1 s−1 = 2.409x102 Time = 241 seconds...