Rates reaction drcraig

Rates of Reaction“The iodide Clock”

By: Dr. Robert D. Craig, Ph.D.For Freshman Chemistry 2

Measuring Reaction Rates

• In this experiment we will determine the effect of a reactant concentration and temperature on the rate of a chemical reaction. A reaction is chosen which proceeds conveniently slowly near room temperature and which can be measured easily by a dramatic color change.


• By measuring the initial rate (the rate near reaction time zero) for a series of reactions with varying concentrations, we can deduce to what power the rate depends on the concentration of each reagent.

• Demonstration of variations in reaction rates:

H2O2 vs. I- vs. H30+ in water to produce I2-

H2O2(aq) + 2I-(aq) + 2 H+(aq) ->

I2(aq) + 2 H2O(aq)


• this reaction proceeds, the colorless reactants gradually develop a brown color due to the product I2. Because of the difficulty of timing the appearance of the I2, we make use of another much faster reaction in the same solution to mark the progress of the slow reaction

Fast Reaction

I2 + 2 S2O32- -> 2 I - + S4O6

2-

Fast!!!!


• The reaction happens so fast that I2 produced by reaction 1 is consumed instantaneously by the thiosulfate (S2O3

2-), so that the I2 color cannot develop. Because both S2O3

2- and S4O62- are

colorless, the solution remains colorless. However, we do not add enough thiosulfate to react with all the I2 that will be formed from reaction 1.


• By this device the reaction solution stays colorless until the instant at which all the thiosulfate is consumed, and then free I2 begins to appear.


• We time the reaction from the initial mixing until the appearance of I2.

• In order to help see this appearance we add starch indicator which forms an intensely colored dark complex with I2 and signals the appearance of I2 by a dramatic color change.


• We know the initial concentration of S2O32-

and measure the time interval for this S2O32- to

be consumed. These quantities determine the rate of the slow reaction 1. By repeating the experiment with different concentrations of H2O2 and different temperatures, we can determine the effect of H2O2 concentration and temperature on the rate of reaction 1.


• .

Effect of Acidity• This experiment tests the effect of different

sulfuric acid concentrations on the rate of the iodine clock reaction. The initial concentrations and volumes of all other reactants are the same in all four vials. The greater the acid concentration, the faster the reaction.

• H2SO4 Concentration Ratios (1 :0.83 : 0.67 : 0.51)

Effect of Acidity• .

Effect of Iodide

• This experiment tests the effect of different iodide ion concentrations on the rate of the iodine clock reaction. The initial concentrations and volumes of all other reactants are the same in all four vials. The greater the iodide concentration, the shorter the reaction time.

• Iodide Ion Concentration Ratios (1 : 0.75 : 0.50 : 0.375

Effect of Iodide• .

Effect of Peroxide

• This experiment tests the effect of different hydrogen peroxide concentrations on the rate of the iodine clock reaction. The initial concentrations and volumes of all other reactants are the same in all four vials. The greater the peroxide concentration, the less time it takes for the solution to turn blue.

• H2O2 Concentration Ratios (1 : 0.80 : 0.60 : 0.40)

Effect of Peroxide

• .

Effect of Temperature

This experiment tests the effect of temperature on the rate of the iodine clock reaction. The reaction rate is greater at higher temperatures.

Effect of Temperature

http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/STILLS/CLOCKRX/CLOCK6/64JPG48/0.JPG

Measuring Reaction Rates RUN 1

•

Dispense reagents for Run 1 into another set of glassware.

• Have a clock or watch with a second hand ready. Then in quick succession:

Add the contents of both beakers B and C into flask A.

Swirl the mixture in A.

Measuring Reaction Rates RUN 1

– Start timing as soon as beaker B is emptied. This is time ti .

– • Place a thermometer in flask A but DO NOT STIR WITH IT.

– • Record the time interval Δt = tf – ti in seconds.


• RUN 2 – • Thoroughly rinse flask A just used, and then

dispense reagents. – • Follow the procedure for Run 1. This one is

faster.


• RUN 3 – • Thoroughly rinse flask A, and then dispense

reagents. – • Follow the procedure for Run 1. This is even

faster than 2.

Measuring Reaction Rates• RUN 4

– • Thoroughly rinse flask A, and dispense reagents. – • Place flask A and beaker C (but not B) on a hot plate

and warm these up to between 45 oC and 50 oC. Use a thermometer to monitor the temperature in flask A only.

– • When flask A is over 45 oC, remove both from the hot plate. Mix B and C into A and swirl.

– • Leave the thermometer in the reaction mixture and record the temperature when the color turns. This one is very fast.


RUN 5 • The temperature of solution A in the ice

bath should be below 10 oC.

Measuring Reaction RatesOn the data sheet, fill in the following quantities

in the table for each of the five runs: • Row 1: Δt is your measured time interval in

seconds. • Row 2: Temp, oC is the temperature of the

reaction mixture. • Row 3: [H2O2] is the [H2O2]molarity in flask

A at time t i . • Calculate the total volume ignoring the 2

drops of starch.


• Row 4: [S2O32-]i is the thiosulfate

molarity in flask A at time t i . • Row 5: Δ[H2O2] is the change in

peroxide concentration during Δt. • This calculation is illustrated below.


• Below is a graph plotted following the iodine clock experiment. The graph could be of any reaction where you can measure the concentration of a reactant over time.



1. Set up an ice bath. 2. Dispense reagents into one set of

glassware A, B, and C as shown above for Run 5. Set these in the ice bath for later use.

Examples of factors affecting rate• For each of the 5 factors give one example

from today’s demonstrations• Nature of reactants –

• Ability of the reactants to meet –

• The concentration of the reactants –

• The temperature of the system –

• The presence of catalysts –

Examples of factors affecting rate• Nature of reactants – H202 is very reactive

• Ability of the reactants to meet – Most reactions occur in the aqueous phase

• The concentration of the reactants – How does this effect rate

• The temperature of the system – reactions that are exothermic are faster due to room temperature

• The presence of catalysts – an oxidizing reagent like MnO2 is a catalyst for this reaction

Measuring Reaction RatesAnswer these questions:1. What units are associated with concentration?2. What units are associated with reaction rate?3. What do the square brackets in [H202] indicate 4. Explain how the rate of a reaction is

determined 5. Plot this data (include title, axes labels):

H202 0.100 0.072 0.056 0.046 0.039 0.034 0.030 0.026

Time (s) 0 50 100 150 200 250 300 350


• Rate = k[I-]p [H2O2]q[H3O+]r

• This reaction is the oxidation of iodide ion (I-) to molecular iodine (I2) by hydrogen peroxide (H2O2):

• It is what happens when no thiosulfate is present “A BACK REACTION”


• Because we can measure the concentrations in the rate law using the techniques described above, the unknowns we wish to measure are k, p, and q. One method of directly measuring k, p, and q is called the method of initial rates.


1. How does the rate at the beginning of a reaction compare to the rate later in a reaction?

2. Explain how the rate of reaction of H2O2(aq) + 2I-(aq) + 2 H+(aq) -> I2(aq) + 2 H2O(aq) is determined experimentally.

Measuring Reaction Rates1. Concentration: mol/L or mol•L–1

2. Rate = concentration/time: (mol/L)/s or mol•L–1•s–1

3. The square brackets "[ ]" is the symbol for concentration (mol/L)

4. The rate of reaction is measured by:Instantaneous slope; rise over run; slope of the tangent at any point.


Slope (rise/run) is the reaction rate in (mol/L)/s…

H202 concentration (mol/L) 0.1 0.072 0.056 0.046 0.039 0.034 0.03Time (s) 0 50 100 150 200 250 300

H202 concentration (mol/L) Time (s) Concentration of H2S04 (mol/L) Concentration of KI(mol/L)0.1 0 0.1 0.1

0.072 50 0.072 0.0720.056 100 0.056 0.0560.046 150 0.046 0.0460.039 200 0.039 0.0390.034 250 0.034 0.0340.03 300 0.03 0.03

0.026 350 0.026 0.026

Measuring reaction rates

Measuring Reaction Rates cont1. As the reaction proceeds, the rate decreases

because reactants are being used up (recall, concentration of reactants affects rate)

2. H202 concentration is measured indirectly by measuring the production of I2(g) (purple) – likely via a spectrometer

For more lessons, visit www.chalkbored.com

http://www.chalkbored.com/

Measuring Reaction Rates• Through experiments and calculations, the

rate law of the Iodine Clock Reaction at approximately 24°C was found to be…

• Rate = 1.18 X 10-4 [I-] [H2O2] • Using the rate law, the rate of the reaction

can be found at the given temperature of 24°C with any molar concentration of I- and H2O2.

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Rates reaction drcraig