Hydrolysis Lab Report PDF

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Hydrolysis of t-Butyl Chloride Evan Parker, Nitin Gharpure, Sarah Slavik Group 6 22 October 2019

Introduction

In a hydrolysis reaction, water is used to break down a compound by detaching covalent bonds and replacing them with water molecules. This is apparent in the hydrolysis of t-butyl chloride performed in this lab. The hydrolysis of t-butyl chloride is special in that it is a substitution, nucleophilic, unimolecular reaction. This suggests that the hydrolysis of t-butyl chloride takes place by means of an SN1 reaction mechanism composed of a slow step where chlorine leaves, thereby generating a carbocation intermediate before water participates in nucleophilic attack onto the carbocation and then onto the intermediate formed in fast steps. This mechanism is shown in Figure 1.

Figure 1 (above): SN1 Reaction Mechanism for Hydrolysis of T-Butyl Chloride The chemical kinetics of this reaction can be determined by rate law. Multiple reactions at different temperatures and concentrations were used to determine the rate law of t-butyl chloride under a variety of conditions. A generic rate law involving two species [A] and [B] is expressed in Equation 1. Rate=k [ A] a [ B]b Equation 1 (above): Generic Rate Law The concentrations of reagents used in the reaction are represented by the letters in the brackets, while the letter k represents the rate constant. The superscripts a and b represent the order of the reaction and given the reaction above the reaction it is a first order reaction. It must be noted that a and b refer to experimentally determined constants. The slow step in a reaction serves as a “bottleneck”, and determines the overall rate. As per Figure 1, the slow step in the hydrolysis of t-butyl chloride is the loss of chlorine. As such, the only species present is t-butyl chloride itself and because the reaction mechanism is unimolecular, a rate law for the reaction is proposed below.

Rate = k[ClC(CH3)3]1 Reactions are affected by the conditions where the reaction takes place, such as concentration, the presence of a catalyst, and temperature. Because the rate constant is affected by the temperature of the reaction to relate the two factors the Arrhenius equation is used: ln k1/k2 = -Ea/R(1/T2 – 1/T1) For this equation k1 and k2 represent the rate constants at their relative temperature. Temperature 1 is the temperature at constant 1 and temperature 2 is the temperature at constant 2. R value is a constant that equals 0.008314 kj/(mol*K). Ea represents the activation energy. When plotting an Arrhenius plot using the reactions making 1/T the x axis and ln(k) the y axis, the slope of that line is -Ea/R. Since R is a known constant activation energy can be found by multiplying the slope of the plot by -R. Table 5. Table of Reagents Reagents

MW (g/mol)

MP (°C)

BP (°C)

Density (g/cm3)

T-Butyl Chloride

92.57

-26

51

0.84

NaOH

39.997

318

1388

2.13

Bromophenol Blue Indicator

669.96

273

279

2.2

58.08

-95

56

0.784

18.01

0

1001

1

Acetone

Water

Experimental To begin the experiment, Part C required a hot both and cold a bath. A 400ml beaker with water was placed on a hot plate which was heated to at least 10 degrees Celsius above room temperature. Ice was also added to a 400 ml beaker with water which was cooled to 10-15 degrees Celsius. Part A of the experiment called for 3.0 mL of a 0.10 M t-butyl chloride solution to be placed in a dry 25 ml Erlenmeyer flask. A sheet of white paper was placed underneath to aid in the observation of the reaction of color change as it reached equilibrium. The next step required 0.3 mL of a 0.10 M NaOH solution, 6.7 ml of distilled water and three drops of bromophenol blue indicator were added to a 50 ml Erlenmeyer flask and swirled. A very important note to add is that different pipettes were used to obtain reagents to prevent possible cross contamination and unwanted reactions. The t-butyl chloride solution in the flask was then poured into the 50 ml flask, and a timer on a phone was started. The timer was on to record the amount time it took for the solution to change color. The mixed solution was swirled then placed

backed into 25 ml flask. The flask was swirled continuously until the color change occurred. Once this happened, the timer was stopped and the time was recorded. A thermometer was placed in the solution to take the temperature of the solution directly after the color change. Once both the temperature and time (in seconds) were recorded the both flasks were rinsed with water, rinsed with acetone, and allowed to air dry. This process was repeated for two more trials. Part B the setup for the reaction was the same as part A except, in addition to 0.3 mL of a 0.10 M NaOH solution, 6.7 ml of distilled water and three drops of bromophenol blue, 10 mL of a 70% water and 30% acetone mixture was also added to the 50 ml flask. As with part A, for the two solutions in both flasks were mixed. The temperature and time were recorded at the color change. This was repeated for each of the three total trials, noting that after each trial the both flasks were rinsed with water, rinsed with acetone, and allowed to air dry. Part C had an identical set up process to Part A. However, for each the cold trials before the mixing of the two solutions occurred, both flasks were placed in the cold bath for five minutes. After five minutes the two solutions were mixed in the 50 ml flask, poured back into the 25 ml flask, and placed in the cold bath until the color change occurred. A thermometer was placed in the bath to record the temperature and the time of the color change was also recorded. For the hot bath trials, before each of the solutions were mixed the two flasks were placed in the hot bath for 5 minutes. Then, after five minutes the two solutions were mixed in the 50 ml flask, poured back into the 25 ml flask, and placed in the hot bath until the color change occurred. A thermometer was placed in the hot bath to note the temperature at the time of the color change and timer was used to record the time of color change in seconds. After each of all six trials (3 hot and 3 cold), both flasks were rinsed with water, rinsed with acetone, and allowed to air dry. This resulted in a total of 12 trials, each with a recorded temperature and time if color change. Results For Part A of the experiment, the temperature of the solution and the time in seconds of the color change was recorded. Table 1 below shows this data. Table 1. Part A: Establishing a Baseline Trial

Temperature (°C)

Seconds

Trial 1

23.5

39.3

Trail 2

23.5

30

Trial 3

23.5

43.57

Part B of the experiment, the temperature of the solution and the time in seconds of the color change were recorded. Table 2 below shows this data. Table 2. Part B: Effect of a Change in Concentration Trial

Temperature (°C)

Seconds

Trial 1

23.7

55

Trail 2

23.7

50

Trial 3

23.7

75

For Part C (hot) Table 3 shows recorded temperature and time in seconds of the color change. Table 3. Part C: Temperature Dependence (Hot) Trial

Temperature (°C)

Seconds

Trial 1

35.6

12.6

Trail 2

35.6

10.8

Trial 3

35.6

11.4

For Part C (cold) Table 3 shows recorded temperature and time in seconds of the color change. Table 4. Part C: Temperature Dependence (Cold) Trial

Temperature (°C)

Seconds

Trial 1

13.6

214

Trail 2

11.2

271

Trial 3

12

294

Utilizing the molarity of the reagents the concentrations, which will be needed for the integrated rate law, were calculated below. Part A and C used the same amounts of reagent in their trials so only the calculations are equivalent. Part B had different amounts of reagents, so its calculations of molarity are shown as well.

Molarity = mol of solute / Liters of solution Part A and C: t-Butyl Chloride: 0.1 M = mol/0.003L mol=0.003 0.003/0.01L=0.03M Concentration of t-butyl chloride in final solution = 0.03 M NaOH: 1.1 M = mol/0.0003L mol=0.00003 0.00003/0.01L=0.003M Concentration of NaOH in final solution = 0.003 M

Water: 1 M = mol/0.0067L mol=0.0067 0.0067/0.01L=0.67M Concentration of water in final solution = 0.67 M Part B: t-Butyl Chloride: 0.1 M = mol/0.003L mol=0.0003 0.003/0.02L=0.015 M Concentration of t-Butyl chloride in final solution = 0.015 M NaOH: 0.1 M = mol/0.0003L mol=0.00003 0.0003/0.02L=0.0015M Concentration of NaOH in final solution: 0.0015 M Water: 1 M = mol/0.0137L mol=0.0137 0.0137/0.02L=0.685 M Concentration of water in the final solution: 0.685 M

The rate of the reaction was first order so to the integrated rate law is: ln ([At]/[Ao] = -kt A0 represents the initial concentration and At represents the final concentration. K is the rate constant and t is the time elapsed. The first order rate law for the hydrolysis of t-butyl chloride is: ln ([ClC(CH3)3]t / [ClC(CH3)3]0] = -kt This equation was used to find the k values of each part of the experiment. The average k value of each part was recorded. Because this law is temperature dependent part B calculations were not necessary Calculation of k Part A: k = ln (0.03/0.1) = -k x 39.3 = 0.030635 k = ln (0.03/0.1) = -k x 30 = 0.040132 k = ln (0.03/0.1) = -k x 43.57 = 0.027633 k (average) = 0.0328 Calculation of k Part C (Hot): k = ln (0.03/0.1) = -k x 12.6 = 0.095553 k = ln (0.03/0.1) = -k x 10.8 = 0.111479 k = ln (0.03/0.1) = -k x 11.4= 0.105612 k (average) = 0.10421

Calculation of k Part C (Cold): k = ln (0.03/0.1) = -k x 214 = 0.005626 k = ln (0.03/0.1) = -k x 271 = 0.004443 k = ln (0.03/0.1) = -k x 294 = 0.004095 k (average) = 0.00472

Using the average temperatures and k values from above an Arrhenius plot was made. The ln of each average constant was on the y-axis and the inverse temperature (1/T) in kelvins was put on the x-axis. The slope of the line is equal to -Ea/R. Ea is the activation energy and R is a constant with a value of 0.008314. The slope is shown in the graph below. Activation energy was: 1755=Ea/0.008314 Ea=14.59

Discussion Reagent concentrations influence reaction rate. In the case of a unimolecular reaction like in hydrolysis of t-butyl chloride, the starting slow step involves chloride leaving the molecule. For this to occur, the molecule must have a high enough energy. While the proportion of t-butyl chloride molecules with sufficient enough energy to begin the slow step of the reaction does not change because the concentration is altered rather than the temperature, an increase in concentration will still increase the total number of molecules in an energetically favorable state to begin the reaction. The influence of reagent concentration on reaction rate is apparent in parts A and B of the experiment. Part A had a higher concentration of t-butyl chloride as compared to

Part B. Resultantly, it led to a shorter time until the reaction reached its endpoint – the reaction rate was greater than that in Part B. This observation may be generalized to most reactions, if a reaction solely depends on its reagents in order to begin (and not on say, a catalyst), then increasing reagent concentration allows for an increase in reaction rate and vice versa. The temperature of a reacting solution also affects the rate of a reaction. The rate limiting step in the hydrolysis of t-butyl chloride involves the breaking of a bond between C and Cl. Heating a solution increases its temperature, meaning that the individual kinetic energies of the t —butyl chloride atoms are increased. By adding energy through heat, more t-butyl chloride molecules are able to reach an energetically favorable state for the slow step of the reaction to occur. This was apparent in part C of the experiment. The hot reactions reacted at a significantly higher rate than the cold reactions. The experimentally determined activation energy was 14.59 kJ, a substantially lower value than that in literature, of 79-84 kj/mol . It is important to note that the reagent, t-butyl chloride, is dissolved in acetone, a chemical which evaporates at room temperature. Because the cold reactions took a long time to prepare, some of the acetone may have evaporated, thereby increasing the concentration of t-butyl chloride relative to what was expected. This makes the reaction go faster, and failure to account for the increase in concentration of t-butyl chloride makes one liable to overestimate rate constant for the reaction, which leads to a smaller difference in the impact of temperature and therefore activation energy. It is also important to note that the solutions prepared for the reaction were prepared in batches, with different liquids added directly to the flask. Preparing the solutions ahead of time would ensure uniform solution quality, and less noise in the data. Conclusion The increase in the rate of reaction lowers the activation energy of a reaction. Higher reagent concentration increases the rate of a reaction. Lower temperatures slow the rate of a reaction due to the lower kinetic energy. Due to human errors in the lab, the correct activation energy was not found. Attention to detail would greatly improve the results of the lab such as knowing the chemicals in the pre-mixed solutions. Group 6 dealt with numerous problems such as not being careful with the solutions by placing the beakers on un-even surfaces, causing the mixed solution to spill causing to prepare more solutions numerous times. Group 6 also dealt with incorrectly doing part A as extra acetone was added to the T-Butyl Chloride solution. When attempting the first part of the experiment, the equilibrium was never met due to the acetone taking over the T-Butyl Chloride. This caused it to be too diluted and no reaction occurred. Again, by not being more knowledgeable, this caused a delay of finishing the experiment

References 1. Casselman, B. Hydrolysis of t-Butyl Chloride. UAB’s Organic Chemistry Background for Organic Chemistry Lab. 2. PubChem. (n.d.). Retrieved from https://pubchem.ncbi.nlm.nih.gov/. 3. Casselman, B. Hydrolysis of t-Butyl Chloride. UAB’s Organic Chemistry Procedure for Organic Chemistry Lab. 4. Hunt, I. R. http://www.chem.ucalgary.ca/courses/351/Carey5th/Ch08/ch8-2.html (accessed Oct 19, 2019)....