Online Experiment 5
Rate and Order of a Chemical Reaction: Iodination of Acetone
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Student Name:
Date Performed:
Rate and Order of a Chemical Reaction:
Iodination of Acetone
OBJECTIVES
- To determine the orders of reaction with respect to acetone, and hydrogen ion .
- To confirm that the order of reaction with respect to iodine is zero.
- To evaluate the rate constant k for the mixtures prepared at room temperature.
Lecture Topic References
Review the following before performing the experiment:
Chapter Title: Chemical Kinetics
Section Title: The Rate of a Chemical Reaction; Measuring Reaction Rates; The Rate Law: The Effect of Concentration on Reaction rate
Tro, N. (2017). Chemistry: A Molecular Approach. 4th Edition. Boston: Pearson. (Or latest edition)
CONCEPTS
The rate at which a chemical reaction occurs depends on several factors: the nature of the reaction, the concentrations of the reactants, the temperature, and the presence of possible catalysts. All these factors can markedly influence the observed rate of reaction. Some reactions at a given temperature are very slow, for example, the oxidation of gaseous hydrogen or wood at room temperature would not proceed appreciably in a century. Other reactions are essentially instantaneous. The precipitation of silver chloride when solutions containing silver ions and chloride ions are mixed and the formation of water when acidic and basic solutions are mixed are examples of extremely rapid reactions. In this experiment you will study a reaction that, in the vicinity of room temperature, proceeds at a moderate, relatively easily measured rate.
For a given reaction, the rate typically increases with an increase in the concentration of any reactant. The relation between rate and concentration is a remarkably simple one in many cases, and for the reaction
aA + bB → cC
The rate can usually be expressed by the equation
rate = k(A)m(B)n (1)
Where m and n are generally, but not always, integers, 0, 1, 2, or possibly 3; (A) and (B) are the concentrations of A and B (in moles per liter); and k is a constant, called the rate constant of the reaction, which makes the relation quantitatively correct. The numbers m and n are called the orders of the reaction with respect to A and B. If m is 1, the reaction is said to be first order with respect to the reactant A. If n is 2, the reaction is second order with respect to the reactant B. The overall order is the sum of m and n. In this example the reaction would be third order overall.
In this experiment you will study the kinetics of the reaction between iodine and acetone:
O O
II II
CH3 – C – CH3(aq) + I2(aq) → CH3 – C – CH2I(aq) + H+(aq) + I–(aq)
The rate of this reaction is found to depend on the concentration of hydrogen ion in the solution as well as presumably on the concentrations of the two reactants. By Equation 1, the rate law for this reaction is
rate = k(acetone)m (I2)n (H+)p (2)
Where m, n and p are the orders of the reaction with respect to acetone, iodine, and hydrogen ion, respectively, and k is the rate constant for the reaction.
The rate of this reaction can be expressed as the (small) change in the concentration of I2, ∆(l2), that occurs, divided by the time interval ∆t required for the change:
rate = (3)
The minus sign is to make the rate positive (∆(l2) is negative). Ordinarily, since rate varies as the concentrations of the reactants according to equation 3, in a rate study it would be necessary to measure, directly or indirectly, the concentration of each reactant as a function of time; the rate would typically vary markedly with time, decreasing to very low values as the concentration of at least one reactant becomes very low.
The reaction kinetics of iodination of acetone is one that can be investigated easily. First of all, iodine has color, so that one can readily follow changes in iodine concentration visually. A second and very important characteristic of this reaction is that it turns out to be zero order in l2 concentration. This means (see Equation 2) that the rate of the reaction does not depend on (I2) at all; (I2)0 = 1, no matter what the value of (I2) is, as long as it is not itself zero.
Because the rate of the reaction does not depend on (I2), you can study the rate by simply making I2 the limiting reagent present in a large excess of acetone and H+ ion. You then measure the time required for a known initial concentration of I2 to be used up completely. If both acetone and H+ are present at much higher concentrations than that of I2, their concentrations will not change appreciably during the course of the reaction, and the rate will remain, by Equation 3, effectively constant until all the iodine is gone, at which time the reaction will stop. Under such circumstances, if it takes t seconds for the color of a solution having an initial concentration of I2 equal to (I2)0 to disappear, the rate of the reaction, by equation 3, would be
rate = (4)
Although the rate of the reaction is constant during its course under the conditions they have set up, you can vary it by changing the initial concentrations of acetone and H+ ion. If, for example, you should double the initial concentration of acetone over that in Mixture 1, keeping (H+) and (I2) at the same values they had previously, then the rate of Mixture 2 would, according to Equation 2, be different from that on Mixture 1:
rate 2 = k(2A)m(I2)0(H+)p (5a)
rate 1 = k(A)m(I2)0(H+)p (5b)
Dividing the first equation by the second, we see that the k’s cancel, as do the terms in the iodine and hydrogen ion concentrations, since they have the same value in both reactions, and we obtain simply
(6)
Having measured both rate 2 and rate 1 by Equation 4, we can find their ratio, which must be equal to 2m. We can then solve for m either by inspection or using logarithms and so find the order of the reaction with respect to acetone. For example, in equation 6, the ratio of
rate 2/rate 1 = 2. Calculate for m,
By inspection approach: 2 = 2m
Question: What is the value of m that will give the answer 2?
Answer: m = 1
Use your calculator to check: 21 = 2, therefore, m = 1
By logarithm approach: Convert both sides to logarithmic expression,
log 2 = m log 2
m =
m = 1
By a similar procedure we can measure the order of the reaction with respect to H+ ion concentration and also confirm the fact that the reaction is zero order with respect to I2. Having found the order with respect to each reactant, we can evaluate k, the rate constant for the reaction at room temperature.
The determination of the orders m and p, the confirmation of the fact that n, the order with respect to I2, equals to zero, and the evaluation of the rate constant k for the reaction at room temperature comprise your objectives in this experiment.
MATERIALS
4 M Acetone
0.005M Iodine
1 M HCl
DI water
Two 50-mL Erlenmeyer flasks
White copy paper
Stopwatch (mobile phone can be used)
Foil square to cover your iodine solution
PROCEDURE
Instruction: Obtain your Data Group assignment (Group A, B or C) from your instructor. As you read the procedure, understand the application of various kinetic concepts and data collection techniques (as shown on the video) to determine the rate of reaction of iodination of acetone, and order of reaction with respect to each of the three reactant. Use the data set assigned to you by your instructor shown in the Data Table after the procedure section.
Table 1 Quantities for Rate Law Determination
| Mixture Number | 4 M acetone (CH3COCH3) | 1 M HCl | DI water | 0.005 M Iodine (I2) |
| I | 5 mL | 5 mL | 10 mL | 5 mL |
| II | 10 mL | 5 mL | 5 mL | 5 mL |
| III | 5 mL | 10 mL | 5 mL | 5 mL |
| IV | 5 mL | 5 mL | 12.5 mL | 2.5 mL |
CH3COCH3(aq) + I2(aq) CH3COCH2I (aq) + H+ (aq) + I– (aq)
- Use the quantities in Table 1 to study the kinetics of the above reaction.
Note: The concentrations given in the table are the concentrations of stock solutions, not the initial concentration of reactants in each mixture. Calculate the initial concentrations ( of acetone, hydrochloric acid and iodine solutions using the equation below. Record in the Data and Calculations section.
(Equation 1)
- Using the quantities in Table 1, prepare Mixture I by adding acetone, HCl and DI water (in this order) inside a 50-mL Erlenmeyer flask.
- Before adding the Iodine solution last in the same Erlenmeyer flask (will be referred to as “Reaction Flask”), make sure your stop watch is ready. It is also necessary to prepare a “Reference Flask” by adding 25 mL of DI water in a clean 50-mL Erlenmeyer flask. Place both flasks on a white copy paper.
Note: Apart from Iodine (I2) being the limiting reactant (acetone and hydrochloric acid are in excess amounts), it is yellow brown in color. Hence, its color change in the solution will serve as the indicator of reaction completion. It is, therefore, important to add the iodine solution last in each reaction. Since iodine solution is light sensitive, cover it with a foil after taking a sample from the reagent bottle and while waiting for it to be added in your work bench.
- While noting the time on your stopwatch, add the indicated amount of iodine solution into the Reaction Flask from Step 2. (Note: Starting the stopwatch and adding the iodine solution into the Reaction Flask must be done simultaneously). Carefully swirl the flask and continue to do so until the color completely disappears. Use the while paper as a background. Record the time it takes for the iodine color to disappear. (Compare the appearance of Reaction Flask content with that of Reference Flask against the white copy paper background).
Note: Use the Reaction Flask containing Mixture 1 as the Reference Flask for Mixtures II to IV
rate determinations. - Measure the temperature of the content inside the Reaction Flask. Record in the Data Sheet.
- Repeat Steps 2 to 5 to determine the reaction rates for Mixtures II to IV.
- Calculate the reaction rates for each mixture, using the initial concentration of I2 (mole/L or M) and reaction time (seconds).
Reaction Rate = (Equation 2) - Using the general rate law expression below, and the calculated rates and initial concentration values, write the complete rate law expression (Rate I to Rate IV) for each mixture.
Rate = k [acetone]m [I2]n [H+]p
- Compare the rates and determine the ratio between these rates. Check if the trend coincides with the change in concentration of any of the reactants (for example, the reaction rate doubles when acetone initial concentration doubles). Using this trend, select the two rate laws that correspond to these parallel changes to calculate the order of reaction with respect to the target reactant.
- Get the ratio of the rate law with the higher rate and the rate law with a lower rate (divisor). Determine the correct ratio of the two rate laws to calculate the reaction order with respect to acetone (m), iodine (n) and H+ (p) as per the general rate law equation in Step 7.
- Write the complete rate law for the iodination of acetone. Calculate the overall order of reaction and the Rate Constant (k). Express all answers with the correct units.
PROCEDURE VIDEO
Click on the YouTube link https://youtu.be/Lk_o_R1IvVA to see a video of the iodination of acetone experiment.
Data
Below is the summary of the data gathered after performing the Procedure:
Data Table 1 Group A – Iodination of Acetone: Time to Complete Reaction
| Temperature °C |
Time It Took for Mixture Color to Disappear (seconds) | |
| Mixture 1 (D.1) | 23.5 | 201.04 |
| Mixture 2 (D.2) | 23.5 | 101.03 |
| Mixture 3 (D.3) | 23.5 | 99.12 |
| Mixture 4 (D.4) | 23.5 | 96.06 |
Data Table 1 Group B – Iodination of Acetone: Time to Complete Reaction
| Temperature °C |
Time It Took for Mixture Color to Disappear (seconds) | |
| Mixture 1 (D.1) | 22.5 | 224.33 |
| Mixture 2 (D.2) | 22.5 | 123.15 |
| Mixture 3 (D.3) | 22.5 | 117.43 |
| Mixture 4 (D.4) | 22.5 | 110.01 |
Data Table 1 Group C – Iodination of Acetone: Time to Complete Reaction
| Temperature °C |
Time It Took for Mixture Color to Disappear (seconds) | |
| Mixture 1 (D.1) | 25 | 71.92 |
| Mixture 2 (D.2) | 25 | 35.10 |
| Mixture 3 (D.3) | 25 | 39.11 |
| Mixture 4 (D.4) | 25 | 42.54 |
PROCESSING THE DATA
- Calculate the initial concentrations of acetone, iodine and hydrochloric acid (Use Equation 1 from the procedure).
- Solve for initial rates (Use Equation 2).
- Calculate the order of reaction, m, n, and p, with respect to acetone, H+ and iodine concentrations. Use the approaches mentioned towards the end of the Concepts section.
- Evaluate the rate constants at room temperature.
LAB SAFETY AND WASTE DISPOSAL
Waste Disposal:
Collect and dispose of wastes from Mixtures I to IV, and any unused acetone in the organic waste bottle. Unused iodine and hydrochloric acid solutions must be disposed of in the inorganic waste bottle. All waste bottles are located in the Satellite Accumulation Area (SAA).
Lab Safety:
Wear the appropriate Personal Protected Equipment (PPE). Read all Safety Data Sheets (SDSs) provided by instructor. Pay attention to the safety precautions mentioned in the procedure and by the instructor.
Bibliography
New Jersey City University (2009). Rates of Chemical Reactions, I. The Iodination of Acetone -Slowinski. Signature Labs Series: CHEM 1106/General Chemistry II. Mason, OH: Cengage Learning.
DATA AND CALCULATIONS
Table 1 Reaction Rate Determination
| Mixture Number | Initial Concentrations* | Time for reaction (s) | Temperature (oC)** | Reaction rate (M/s or mole/L ● s) | ||
| Acetone (M) | Iodine (M) | HCl (M) | ||||
| I (D.1) | ||||||
| II (D.2) | ||||||
| III (D.3) | ||||||
| IV (D.4) | ||||||
* Calculate the initial rates before coming to the lab.
** Report the average temperature in the Rate Constant, k, calculation section below.
Show all rate laws/ratios and calculations:
Rate Law Expressions for Mixtures I to IV:
Rate I:
Rate II:
Rate III:
Rate IV:
Rate Ratios (Show all Calculations below): Reaction Order
m____
p ____
n_____
Overall Order of the Reaction: __________
Reaction Rate Law (Rate Equation) for the Iodination of Acetone:
Rate Constant, k :
Mixture I II III IV Average
k ____________ ____________ ____________ ___________ ___________ at _____ oC
Show all calculations:
post-lab questions Score: ________
Using the lessons learned from the concepts, video presentations, and data result analysis, give your best and concise answers to the following questions.
- a) What is the rate of a chemical reaction?
b) What are the factors that affect the reaction rate and which one of these factors did you study
in this experiment?Briefly answer the above questions using the iodination of acetone reaction you performed in this experiment.
- What was the role of iodine, acetone, and hydrochloric acid in the iodination reaction?
- a) How many reactions were performed and which reactant concentrations did you vary in these
reactions?
b) Why did you vary these concentrations? - How was the order of reaction with respect to each reactant determined using the initial rates method? Briefly explain with calculations.
- What is the order of reaction with respect to the concentration of each reactant (acetone, hydrochloric acid and iodine)?
6. Based on your data and calculations, what is your iodination of acetone Rate Law (Rate Equation)? What do these orders of reaction with respect to each reactant indicate with respect to the overall rate of the iodination reaction?
- Using your given data set for mixture 4 and result of your data and calculations section, you were asked to determine the rate of a fifth reaction mixture made up in the following way:
5 mL 4 M acetone + 5 mL 1 M HCl + 15 mL 0.005 M Iodine
How long (in seconds) will it take for the Iodine color to disappear at the same temperature of your Mixture 4?
What practical applications of chemical reaction rate did you learn in this experiment? Give one specific example. (If you use a reference, cite it below using the APA 6th edition format.)
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