Lab 2: Factors Affecting Reaction Rates
Objectives:
To determine how concentration, temperature, and catalyst affect the speed of a chemical reaction.
Introduction:
The rate law is a mathematical expression showing how the rate of the reacting species changes with concentrations. In this experiment, you will determine the rate law of the following reaction by observing how concentration and temperature affect the speed of the reaction.
S2O82- (aq) + 2 I- (aq) 2 SO42- (aq) + I2 (aq) (2-1)
persulfateion / iodide
ion / sulfate
ion / iodine
The factors, which affect the rate of this reaction are:
1. Concentrations of the reactants, S2O82- and I- ions (Part A)
2. Temperature (Part B)
- Catalyst (Part C)
Introduction:
Part A: Effect of Concentration on the Rate of a Reaction
Our method for measuring the rate of the reaction involves what is called a "clock" reaction. The rate law for this reaction is of the form:
rate = k [ S2O82- ]x [ I- ]y eqn (2-2)
where [ S2O82- ] and [ I- ] are the concentration in moles/L of persulfate ion and iodide ion respectively, and k is the rate constant. The rate constant will have a unique value for any particular process at a given temperature. The magnitude of the rate constant will tell whether a reaction will proceed quickly. A small rate constant indicates a slow reaction. A large rate constant indicates a rapid reaction. The rate constant of a chemical reaction will vary with temperature.
Table (2-1) gives of some reactions and their rate equations. The order of a reaction is very useful because it allows us to predict the influence of concentration of the speed of the reaction. For a 1st order reaction (example (c)) doubling the concentration, doubles the reaction rate. But if a reaction is 2nd order (example (a)) doubling the concentration increases the reaction rate by 4 times.
REACTION / RATE LAW / OVERALL ORDER(a) / 2 HI H2 + I2 / Rate = k[HI]2 / 2
(b) / 2 NO + H2 N2O + H2O / Rate = k[NO]2 [H2]1 / 3
(c) / 2 N2O5 4 NO2 + O2 / Rate = k [ N2O5 ] / 1
Table (2-1) - Some chemical reactions and their rate equations.
The exponent 'x' in eqn (2-2) is the order of the reaction with respect to the S2O82- ion. The exponent 'y' is the order of reaction with respect to the I- ion. The overall order of the reaction is the sum of x+y. The powers to which the concentrations are raised, ‘x’ and ‘y’, MAY or MAY NOT be the same as the stoichiometric coefficients in the balanced equation (2-1). In general, the order of a reaction CANNOT be determined by inspecting the balanced chemical equation but must be determined experimentally.
In this experiment, a known amount of sodium thiosulfate, Na2S2O3, is added to the reaction mixture. As reaction (2-1) proceeds, it will start to form I2 (aq). The iodine formed will be consumed according to the following reaction.
I2 (aq) + 2 S2O32- (aq) ® 2 I- (aq) + S4O62- (aq) (2-1-2)
After the S2O32- ions are exhausted, the formation of any more I2 will react with the starch to turn the solution blue.
Part B: Effect of Temperature on the Rate of a Reaction
For almost all reactions, an increase in temperature will lead to an increase in the reaction rate. An increase in temperature will increase the average kinetic energy of the molecules. This will lead to an increase in the number of collisions per unit time. But collision alone is an insufficient criterion for the production of the products. The kinetic energy of the colliding molecules must be greater than the required minimum energy level before the reactants can be converted to products. This energy level is called the activation energy, Eact. Activation energy is unique for a particular chemical reaction.
If a reaction is investigated at a number of different temperatures, the rate constant, k, usually shows quite a dramatic increase, often several orders of magnitude, over a relatively small temperature range. For many reactions there is approximately a two- to three-fold increase in rate for every 10oC rise in temperature. For a given temperature, reactions that have large activation energies would be slower than the ones that have smaller activation energies.
The relationship between the temperature and the rate constant ‘k’ is given by the Arrhenius equation,
eqn (2-3)
or
eqn (2-4)
where
A / is the pre-exponential or frequency factor, a constant related to the collision frequencyR / is the gas constant (8.314 J / K mole)
T / is the absolute temperature (K)
k / is the rate constant at temperature T
Eact / is the activation energy, the energy required by the reacting species for their collisions to be effective (ie - those that lead to the formation of products)
Equation (2-4) shows that 'ln k' is a linear function of the reciprocal absolute temperature. The activation energy can be determined experimentally by measuring the rate constant, k, at several different temperatures, T. A garph of ln k vs. 1/T yields a straight line with a slope of (– Eact / R) and an intercept of ln A.
Part C: Effect of a Catalyst on the Rate of a Reaction
A catalyst is a substance which increases the rate of a reaction, but remains unchanged in the process. A catalyst may function by providing an alternate path for which the reactants come together. In order for the rate of of the catalyst-assisted reaction to increase, the alternate pathway will have a lower activation energy, Eact, and therefore proceeds more rapidly.
In this experiment, a small quantity of Cu2+ is introduced to the persulfate-iodide clock reaction as a catalyst.
Cu2+S2O82- (aq) + 2 I- (aq) 2 SO42- (aq) + I2 (aq) / (2-5)
Apparatus:
1. Water baths set at: 20oC, 30oC, 40oC
2. lead donuts
3. 22 - 125 mL Erlenmeyer flasks (14 for Part A, 6 for Part B, 2 for Part C)
4. 2 plastic buckets to make ice slurry
5. ice
6. Timers that display seconds
7. Alcohol thermometers
Solutions:
1. 0.00500 M Na2S2O3 – in dispenser
2. 0.200 M KI – in dispenser
3. 0.100 M (NH4)2S2O8 – in dispenser
4. 0.1 M CuSO4
- starch indicator
Procedure:
Part A - Effect of Concentration on Reaction Rate
Note: / 1. All Erlenmeyer flasks provided are CLEAN and DRY. Use them as is.2. Part A is carried out at 20oC
1. The following solutions are available in bottle-top dispensers. Pay attention to the preset volumes and dispense the proper amounts into the Erlenmeyer flasks.
(ii) / 0.200 M KI
(iii) / 0.100 M (NH4)2S2O8
(iv) / distilled H2O
2. Obtain fourteen 125 mL Erlenmeyer flasks. Label themas follows:
B1, B2, B3, B4, B5, B6 and B7
3. Prepare the seven ‘A’ solutions according to Table (2-2).
Table (2-2) - Contents of the seven 'A' solutions.
Na2S2O3
(mL) / 0.200 M
KI
(mL) / distilled
H2O
(mL) / 3 % starch
indicator
A1 / 10.0 / 20.0 / 0.0 / 3 drops
A2 / 10.0 / 20.0 / 0.0 / 3 drops
A3 / 10.0 / 20.0 / 0.0 / 3 drops
A4 / 10.0 / 20.0 / 0.0 / 3 drops
A5 / 10.0 / 15.0 / 5.0 / 3 drops
A6 / 10.0 / 10.0 / 10.0 / 3 drops
A7 / 10.0 / 5.0 / 15.0 / 3 drops
4. Prepare the seven ‘B’ solutions according to Table (2-3).
Table (2-3) – Contents of the seven ‘B’ solutions.
(NH4)2S2O8
(mL) / distilled
H2O
(mL)
B1 / 5.0 / 15.0
B2 / 10.0 / 10.0
B3 / 15.0 / 5.0
B4 / 20.0 / 0.0
B5 / 20.0 / 0.0
B6 / 20.0 / 0.0
B7 / 20.0 / 0.0
5.Use lead donuts to stabilize the fourteen Erlenmeyer flasks in the 20oC water bath. Allow the flasks to come to thermal equilibrium by leaving them in the bath for at least 5 minutes. Measure the temperature of the water bath with a thermometer. Record the actual temperatures on the data sheet.
6. Pour the content of solution A1 rapidly into B1 while swirling. Start the timer immediately and record the time (in seconds) for the appearance of the blue colour. For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the 20oC water bath.
7. Repeat step 6 for the remaining 6 pairs of solutions, pouring A2 into B2, A3 into B3, A4 into B4, A5 into B5 and A6 into B6. Record the time (in seconds) for the appearance of the blue colour.
Procedure:
Part B - Effect of Temperature on Reaction Rate
2. Part B is carried out at 0oC, 30oC, 40oC
3. Three sets of A4/B4 solutions will be used at these temperatures.
1. Obtain six 125 mL Erlenmeyer flasks.
2. Prepare three 'A4' solutions and three 'B4' solutions by following instructions given in Table (2-2) and Table (2-3).
3. Obtain two plastic buckets and prepare an ice slurry in each bucket. Use lead donuts to stabilize the Erlenmeyer flasks. Put one pair of 'A4/B4' solutions in the 0oC ice bath. Put the second pair of solution in the 30oC water bath, and the third pair of solution in the 40oC water bath. Allow the flasks to immerse in the bath for 5 minutes to come to thermal equilibrium.
4. Measure the temperatures of the ice slurry and water baths with a thermometer. Record the actual temperatures on the data sheet.
5. Pour the content of solution A4 rapidly into B4 while swirling. Start the timer immediately and record the time (in seconds) for the appearance of the blue colour. For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the water bath.
6. Repeat step 5 for the remaining two pairs of solutions.
Procedure:
Part C - Effect of a Catalyst on Reaction Rate
Note: / 1. All Erlenmeyer flasks provided are CLEAN and DRY. Use them as is.1. Obtain two 125 mL Erlenmeyer flasks.
2. Prepare one 'A4' solutions and one 'B4' solutions by following instructions given in
Table (2-2) and Table (2-3).
3. To solution A4 add 1 drop of 0.1 M CuSO4·5H2O solution.
4. Use lead donuts to stabilize the two Erlenmeyer flasks in the 20oC water bath. Ensure that the temperature of the water bath is the same as the temperature used in Part A. Allow the flasks to come to thermal equilibrium by leaving them in the bath for at least 5 minutes. Measure the temperature of the water bath with a thermometer. Record the actual temperatures on the data sheet.
5. Pour the content of solution A4 rapidly into B4 while swirling. Start the timer immediately and record the time (in seconds) for the appearance of the blue colour. For the duration of the reaction, keep swirling the flask containing the combined solutions and keep the flask immersed in the 20oC water bath.
Datasheet:
Part A - Effect of Concentration on Reaction Rate
Temperature: ______
1 / A1/B1
2 / A2/B2
3 / A3/B3
4 / A4/B4
5 / A5/B5
6 / A6/B6
7 / A7/B7
Datasheet:
Part B - Effect of Temperature on Reaction Rate
(sec) / Temperature
(oC)
4
(Room Temperature) / A4/B4 /
(data copied
from Part A) /
(data copied
from Part A)
8
(in ice) / A4/B4
9
(near 30oC) / A4/B4
10
(near 40oC) / A4/B4
Datasheet:
Part C - Effect of a Catalyst on Reaction Rate
Temperature: ______
11 / A4/B4
Postlab Questions:
Part A - Effect of Concentration on Reaction Rate
1. Transfer data Experiments 1, 2, 3, and 4 from the datasheet to Table (2-4) below and calculate:
· the initial concentration of [S2O82-] and [I -],
· the concentration of I2 formed to consume the S2O32- added,
· the rate of formation of I2.
(a) Show a sample calculation for Experiment 1 in the space provided. Enter the calculated results for Experiment 1 into the first row of Table (2-4) below.
2. Calculate the initial concentration of [I -] for Experiment 1.
[Note: The concentration of [I -] is the same for Experiments 1, 2, 3, and 4]
3. (a) Calculate the concentration of I2 formed to consume the S2O32- added.
(b) Calculate the rate of formation of I2 for Experiment 1.
(b) Repeat the above calculations for Experiments 2, 3, and 4 and complete Table (2-4).
[Note: Data and calculated results for Experiment 4 (highlighted in grey) will be used in other parts of the lab.]
Table (2-4): Summary of calculations for Experiments 1 to 4.
Temperature of Experiments: ______Experiment / Solution / Time
(sec) / [S2O82-]
(moles/L) / [I-]
(moles/L) / Rate of Formation of I2
(M/sec)
1 / A1/B1
2 / A2/B2
3 / A3/B3
4
(Transfer to Table (2-6)) / A4/B4
2. (a) Determine the order of the reaction with respect to [ S2O82- ] using the calculated results from Experiments 1 and 2 from Table (2-4).
[Note: The order of the reaction with respect to [ S2O82- ] is the value of 'x' in the rate equation
rate = k [ S2O82- ]x [ I- ]y .]
Show a sample calculation in the space provided. Enter your calculated result in the first row of Table (2-5).
(b) Repeat the calculation of 'x', using the other pairs of experiments as indicated in Table (2-5) below. Calculate the average value of 'x' . The average value of 'x' should be rounded the nearest integer. Complete Table (2-5).