Lab Report 5 Example - Grade: A PDF

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A Kinetic Study of the Hydrolysis of t-Butyl Chloride

Lead Author: Lindsey Cross Reviewer: Shyla Hossain Editor: Brinda Shah

Introduction

Chemical kinetics is defined as the ‘study of the speeds of reactions and the nanoscale pathways or rearrangements by which atoms and molecules are transformed from reactants to products’.1 Said reactions are labeled bidirectional, meaning they continue in one of two opposite directions until equilibrium is achieved. At that point, the concentration of reactants and products stops changing significantly.2 The rate at which these reactions occur is dependent on the concentration of reactants, temperature of the reaction and solvent polarity, among other factors.3 The reaction rate can be determined by measuring the change in concentration of the reactant divided by the change in time. Eq. 11

Reaction Rate =

Experimentally, all variables are held constant and the only rate determining factor is concentration. This is expressed as the rate law equation.2 Eq. 22 In this equation, k is the rate constant, x is the order of reaction of A, and y is the order of reaction of B.2 The order of reaction describes the effect of reactant or product concentration on reaction rate and can be first, second, or third order. The overall order of reaction is determined by adding x and y. This order can only be determined experimentally.1 The current experiment studies the effect of reactant concentration and temperature on the hydrolysis of t-Butyl chloride. It is a nucleophilic substitution reaction. This reaction has two limiting reactions SN1 and SN2. The SN1 mechanism

means that the bond breaking is completely done before any bond formation can begin. The SN2 mechanism means that bond breaking and forming happen at the same time.4 The hydrolysis of t-Buytl is a SN1 reaction. The rate-determining step cleaves a chlorine ion from a carbon atom, forming an intermediate carbocation. Then, the nucleophile, water, quickly reacts with the carbon to satisfy the carbocation, by forming a bond. In the last step, a hydrogen ion cleaves from the compound to form hydrophilic acid and converting the alkyl halide to an alcohol as seen in Figure 1.1

Figure 1 Mechanism of the hydrolysis of t-Butyl chloride

Table 1 All reagents used in the experiment. Name

Molecular Weight (g/ mol)

Melting Point (°C)

Boiling Point (°C)

Density (g/cm3)

Hydrochloric Acid 36.46

-114.2

-85.1

1.49

Water

0.0

100

1.0

Sodium Hydroxide 40.0

318.0

1,388

2.13

Bromphenol Blue

670

279

206

2.2

Acetone

58.1

-94

56.5

0.79

t-Buytl Chloride

92.57

-26.0

51.0

0.84

18.02

t-Buytl Alcohol

74.12

25.5

82.5

0.78

Experimental The procedures necessary to perform the experiment involved many steps and which involved many steps and were consequently separated into parts A, B, and C. In part A, three milliliters of a 0.1M solution of t-butyl chloride in acetone was added to one flask, while7 milliliters of a 10% NaOH and water solution was added to a separate flask with a few drops of bromphenol blue. Next, the contents of both flasks were thoroughly mixed; as the solution settled, we observed and recorded the time it took the coloration of the solution to change form blue to yellow – a clear indication of a pH of 4.6. This procedure was repeated two more times; after the final run all three times were averaged. The procedure for part B was the same for part A except 10mL of a 70/30 water/acetone solution was added to the NaOH mixture. This halved the concentration of the t-butyl chloride. Part C was also the same as part A, except three runs were done at 15°C and three runs were done at 35°C. Results The raw data and average colorization times for Parts A, B, and C are shown in Table 2.

Table 1. Raw and average reaction times for all parts of the experiment. Runs

Part A (sec)

Part B (sec)

Part C (15°C) (sec)

Part C (35°C) (sec)

1

70

64

70

11

2

56

66

75

12

3

53

84

83

12

Average

60 (9.07)

71 (11.0)

76 (6.56)

12 (0.577)

Solutions in Part A were prepared from 0.1M t-butyl chloride in acetone, mixed with a 10% NaOH and water solution. Part B was prepared with a 70:30 water-acetone solution. Part C kept the concentrations of Part A and changed the temperature. The time was recorded at the point bromphenol blue changed to yellow. Standard deviation is shown in parentheses.

In order to solve for the rate constant, the molarity must be calculated. Equation three shows the conversion of molarity to moles for t-butyl chloride and sodium hydroxide. M x L = mols

Eq. 3

For Part A and C, the equation is 0.1M t-butyl chloride x 0.003L t-butyl chloride = 3.0x10-4 moles t-butyl chloride. Eq. 4 Calculating the moles of NaOH is done similarly, via: 0.1M NaOH x 0.0003L NaOH = 3.0x10-5 moles NaOH

Eq. 5

The equation for calculating the new molarity of the reactant in solution is Mols/L = M.

Eq. 6

The total volume of the t-butyl chloride and NaOH solution was 10mL for Part A and C for the experiment with 3mL being t-butyl chloride. Consequently, 3.0x10-4 moles of t-butyl chloride/0.01L = 0.003M t-butyl chloride The total volume of NaOH is 3.0x10-5 moles of NaOH/0.01L = 0.003M NaOH. In Part B, the solvent was doubled and the reactants reduced by one-halve, resulting in a ½ decrease in molarity of t-butyl and NaOH. 3.0x10-4moles t-butyl chloride/0.02L = 0.015M t-butyl chloride 3.0x10-5moles NaOH/0.02L = .0015M NaOH. The rate limiting step is the first reaction in which the chloride leaves t-butyl chloride to form a tertiary carbocation. The rate constant is dependent on the temperature of the reaction. Reactions in Part A and B were carried out at room temperature (22-25°C). The equation shows that the reaction is first order. Since k is independent of the concentration of the reactants, it can be calculated by the following eqation, using 10% for the percent reaction variable:

Eq. 7 kt = 2.303 log(1/1-[10/100]) kt = .105

The average time for each of the experiments can be plugged into Eq. 7 to calculate the rate constant. For Part A this is calculated as: k(60) = 2.303 log(1/1-[10/100]) k(60) = 0.105379501/60 k = 1.76 x 10-3 sec-1 For Part B in which the concentration effect was studied, the rate constant would be: k(71) = 2.303 log(1/1-[10/100]) k(71) = 0.105379501/71 k = 1.48 x 10-3 sec-1 For Part C in which the temperature dependence was studied, the rate constant would be calculated as follows: At 15°C: k(76) = 2.303 log(1/1-[10/100]) k(76) = 0.105379501/76 k = 1.39 x 10-3 sec-1

At 35°C: k(12) = 2.303 log(1/1-[10/100]) k(12) = 0.105379501/12 k = 8.78 x 10-3 sec-1 Plotting the natural log of the k values with inverse temperature (1/T) should provide a straight line according to the linear form of the Arrhenius equation. The

slope of the line is equal to Ea/R. R, the gas constant used in this equation, is 8.314 J. Plotting the data to obtain the slope will allow the calculation of the activation energy. Thus, our activation energy is: Ea = -8134(8.314x10-3) = I-67.6I= 67.6 kJ/mol

y = -8134.1x + 21.433 R² = 0.8312

F

Arrhenius Plot ln(k) (1/sec)

0 -1.75 -3.5 -5.25 -7 0.0031

0.0032

0.0033 1/T (K)

0.0034

0.0035

Figure 2. The Arrhenius plot was made from the k values of Part A, Part C at 15°C, and Part C at 35°C. The temperature used for the room temperature experiment (Part A) was 25°C. The linear regression line shows a slope of -8134.1 and a y-intercept of 21.433. Discussion The standard rate of this reaction occurred at room temperature. As previously discussed many things that can change this rate. This experiment revealed that increasing the temperature would then increase the rate of reaction. This is due to the faster movement of the molecules at higher temperatures, causing more collisions as the reaction speeds up. At lower temperatures, the molecules slowed down and fewer collisions occurred, which decreased the rate of reaction. In Part B, the concentration of t-butyl was halved. This caused fewer collisions with NaOH since there were less t-butyl molecules relative to the increased volume1. According to the k values of the 15°C experiment and the concentration experiment, decreasing the concentration seems to decrease the rate of reaction more so than decreasing the temperature. There was possible error that could have affected the data. Different lab partners performed all three parts. The change in technique or mixing between the three could have cause slightly skewed data. Three of the twelve runs performed may have skewed the data as outliers, considering the size of the standard deviations (Table 2). To improve the experiment, many more trials would need to be completed to give a more accurate data set. Conclusion

This experiment and subsequent results demonstrated the effect of concentration and temperature in the rate of reaction on the hydrolysis of t-butyl chloride. Increasing the temperature increased the hydrolysis relative to the room temperature rate of reaction. Additionally, decreasing both the temperature and the concentration of t-butyl slowed the hydrolysis down; lowering the concentration, in particular, exhibited a more pronounced hydrolysis reduction rate. The activation energy was calculated through the use of an Arrhenius plot. Allotting more time or collecting more data points for each experiment could have improved our experiment. As a result, both the limited data and outliers have caused the data to be less accurate than we anticipated.

References (1) Moore, J.; Stanitski, C.; Jurs, P. Chemistry: The Molecular Science; Brooks/Cole: Belmont, CA, 2005. (2) Harris, G. Chemical Kinetics; D.C. Heath and Company: Boston, MA, 1966. (3) Hill, R.; Barbaro, J. Experiments In Organic Chemistry; 3rd ed.; Contemporary Pub. Co. of Raleigh: Raleigh, NC, 2005.

(4) Brown, W.; Foote, C.; Iverson, B.; Anslyn, E.; Novak, B. Organic Chemistry; 4th ed.; Brooks/Cole: Belmont, CA, 2012....