CHEMISTRY

Factors Affecting Reaction Rates

Investigation Manual

FACTORS AFFECTING REACTION RATES

Overview In this activity, the factors affecting the rate of an iodine clock reaction will be investigated. In this reaction, two colorless solutions are combined and the reaction is completed when the solution turns a dark blue color. The reaction rate will be studied by testing the effects of temperature, concentration, and addition of a catalyst on the rate. The measured data will be used to calculate the rate law for the iodine clock reaction.

Outcomes • Determine the effect of concentration, temperature, and catalyst

on reaction rate. • Extrapolate the rate law expression from experimental data.

Time Requirements Preparation ...................................................................... 5 minutes Activity 1: Calibration Tests .......................................... 30 minutes Activity 2: Concentration ............................................... 60 minutes Activity 3: Temperature .................................................. 20 minutes Activity 4: Catalyst ......................................................... 20 minutes Activity 5: Determining the Rate Law ............................ 60 minutes The laboratory experiment may be stopped after the completion of any activity and resumed at a later time. If the experiment is interrupted before it is complete, ensure that all liquids are sealed to prevent evaporation and spillage. Clean the 24-well plate and test tubes and allow them to dry before resuming the next activity.

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Key Personal protective equipment (PPE)

goggles gloves apron

warning corrosion flammable toxic environment health hazard

follow photograph stopwatch link to results and required video submit

Table of Contents

2 Overview 2 Outcomes 2 Time Requirements 3 Background 7 Materials 8 Safety 8 Preparation 9 Activity 1 9 Activity 2 11 Activity 3 12 Activity 4 13 Disposal and Cleanup 13 Activity 5 14 Data Tables

Made ADA compliant by NetCentric Technologies using the CommonLook® software

Background Chemical kinetics seeks to understand the factors that control reaction rates. The funda- mental concept of chemical kinetics is collision theory. According to this theory, reactant mole- cules must collide with enough force to achieve activation energy before a reaction can occur. Activation energy is the minimum amount of energy required to break the chemical bonds of reactant molecules, after which they can form new products. Kinetic molecular theory indicates that molecules continuously collide; however, only those collisions that meet or exceed the required activation energy, and those that collide in a favorable orientation, will form new chemical bonds, thus new products. The rate of chemical reactions can be influenced by temperature, reactant concentration, physical state of the reactants, and the presence of a catalyst.

Temperature Kinetic molecular theory defines temperature as a measurement of the average kinetic energy of molecules. The Kelvin temperature scale begins with 0 K, the temperature at which there is no molecular motion. As temperature increases, the average kinetic energy of molecules also increases, which increases their average speed. At higher speeds, molecules collide more frequently and with higher energy, leading to faster reaction rates. Chemical reactions have faster reaction rates at higher temperatures.

Concentration Concentration is the quantity of molecules per volume. As an example for liquid solutions, 1 Molar hydrochloric acid (1 M HCl) contains 1 mole of HCl molecules per liter of solution,

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whereas 2 M HCl contains 2 moles of HCl molecules per liter of solution. The concen- tration of gas molecules in a closed system is increased by adding more gas per volume, which increases the pressure, or by decreasing the volume of the gas. As molecular concentra- tion of the reactants increases, the frequency of collisions increases, which leads to faster reac- tion rates. Most chemical reactions have faster reaction rates at higher concentrations.

Physical State The physical state of the reactants is important for a reaction to occur. Two solids will have minimal surface area touching one another and hence, very little interaction between the mole- cules of the two solids. The reaction will usually be very slow. If a solid is added to a solution, a reaction may occur at the surface of the solid. However, if the solid is broken up into smaller pieces, the surface area increases and the rate of reaction will also increase.

Reaction rates are affected by surface area because the greater the surface area, the greater the number of molecular collisions. Molecules must collide with each other for a reaction to occur; therefore, the greater the contact surface area, the greater the probability for molecular collisions. Chemical reactions between mole- cules in different phases (e.g., gas-solid or liquid-solid) have slow rates because collisions are limited. By contrast, reactions between molecules in a homogeneous solution have a greater probability for molecular interactions and collisions, and have faster reaction rates.

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FACTORS AFFECTING REACTION RATES

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Step 1: Write the chemical equation:

A + B → Products

Step 2: The reaction rate is written as being proportional to the starting concentration of each reactant. Brackets ([A] and [B]) represent the starting molar concentrations of each reac- tant in moles/L or M. However, each reactant may influence the reaction rate differently. Reac- tion orders are used to indicate the importance of a reactant to the rate determining step, and are represented by the exponents n and m. If the exponent for reactant A is 1, then the reaction is said to be first-order for A. If the exponent is 2, then the reaction is said to be second-order for A. For most reactions, the exponent is 0, 1, or 2. The exponent value must be determined experimentally because it is the extent to which the concentration of a species affects the rate of a reaction. The reaction order cannot be determined from the chemical equation. The proportional rate law is written using the reaction rate, the concentration of each reactant, and the reaction order for each reactant:

R a [A]n[B]m

Step 3: To write the final rate law, the reaction must be run multiple times. The concentration of each reactant is changed in an orderly approach to determine the impact the rate of that reactant has on the reaction rate. Once the order of reac- tion for each reactant is determined experimen- tally, the rate equation is solved by inserting the reaction rate constant (k).

R = k[A]n[B]m

Catalyst A catalyst is a substance that increases the rate of a chemical reaction but is not chemically changed during the reaction. A catalyst provides a mechanism that makes it easier for bonds to break and reform into products. The catalyst lowers the reaction activation energy and greatly accelerates the reaction. For example, the decomposition of H2O2 in an aqueous solution to O2 and H2O is greatly accelerated by the addition of a catalyst. A bottle of 3% hydrogen peroxide will last for years before it completely decom- poses to water and oxygen. If a drop of potas- sium iodide (a catalyst for this reaction) is added to the solution, it will start to react, and within minutes the reaction will be complete. Enzymes are nature’s catalysts and speed up any crucial biochemical reactions in most organisms.

Rate Laws To properly study the rate of a reaction, a rate law must be written. Many reactions occur in multiple steps, and the overall rate of the reaction will depend on the step that takes the longest amount of time. This step is called the rate determining step. The rate law for a reac- tion can only be determined experimentally and provides understanding of the reaction mecha- nism.

The rate of a chemical reaction, also called reaction rate (R) is how long it takes for the reac- tion to occur. It is defined as the concentration change of reactants that occurs during a spec- ified time period. The rate law for a chemical reaction is a mathematical relationship between reaction rate and the concentration of each reactant. Determining the rate law for a chemical reaction can be best explained with the following example and steps:

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The units for reaction rates are usually expressed as molarity per second (M/s) or the change of concentration over time.

The best method to determine the exponents or the reaction order is to vary the concentration of one of the reactants while keeping all other reactant concentrations and reaction conditions constant. For example, examine the following data table for the reaction A + B → Products. In each trial, the starting concentration of one reac- tant is changed while the other reactant concen- tration is held constant, and the reaction time is measured and recorded.

Trial [A] [B] Time(T, sec) Rate (R) = 1/T

(M/s)

1 0.10 M 0.10 M 25.0 0.04

2 0.20M 0.10 M 12.5 0.08

3 0.10 M 0.20 M 6.25 0.16

The results from Trials 1 and 2 show that the reaction rate doubles (from 0.04 to 0.08 M/s) as the concentration of A doubles (from 0.10 to 0.20 M). Consider what the exponent n would have to equal for [A] to satisfy this condition.

[A]n = R [2]n = 2

The exponent n would have to equal 1, indicating that the reaction is first-order for A.

Now consider Trials 1 and 3, which show that the reaction rate increases by a factor of four (from 0.04 to 0.16 M/s) as the concentration of B doubles (from 0.10 to 0.20 M). Consider what

the exponent m would have to equal for [B] to satisfy this condition.

[B]m = R [2]m = 4

The exponent m would equal 2, indicating that the reaction is second-order for B.

If the answer is not easy to determine by simple inspection of the data, the following equation can be used to determine an unknown exponent:

[A2] rate2 [A2] rate1 Using the values given in the preceding data table, solve the equation for n and m, and confirm that the same values are obtained.

To solve for the exponent (reaction order) of [A],

0.20 M 0.08 M/s 0.10 M 0.04 M/s

2n = 2 ln 2n = ln 2

n(ln 2) = ln 2 n = 1

To solve for the exponent (reaction order) of [B],

0.20 M 0.16 M/s 0.10 M 0.04 M/s

2m = 4 ln 2m = ln 4

m(ln 2) = ln 4 In 4

m = In 2

m = 2 continued on next page

( ) n =

( ) m =

( ) m =

FACTORS AFFECTING REACTION RATES

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Trial 3 R = k[A][B]2 k = R / [A][B]2 k = (0.160 M/s) / [0.10 M][0.20 M]2

k = (0.160 M/s) / [0.10 M][0.04 M2] k = (0.160 M/s) / [0.004 M3] k = 40 M-2 s-1

Iodine Clock Reaction The iodine clock reaction performed in this labo- ratory investigation occurs in the following three steps:

Step 1: 3I– (aq) + H2O2 (aq) + 2H+ (aq) → I3 (aq) + 2H2O (l)

Step 2: I3 (aq) + 2S2O3- (aq) → 3I

– (aq) + S4O6– (aq)

Step 3: 2I3 (aq) + starch → starch - I5 complex + l (aq)

2 2

In reaction step 1, iodide ions are oxidized by hydrogen peroxide ions in an acidic solution to form triiodide ions. In step 2, the triiodide ions are reduced by thiosulfate (S2O32–) to form iodide ions. This reaction occurs quickly, and triiodide ions are consumed by thiosulfate as quickly as they are formed. Reaction 3 only occurs when all thiosulfate is consumed. After all the thiosulfate is consumed, triiodide ions are now available to react with starch to form a pentaiodide—starch complex that is a dark blue. The appearance of a dark blue-black color signals the end of the three reactions. The reaction begins when the reactants are mixed together, and ends when the dark blue-black color is observed.

When solving for an unknown reaction order, the only reaction variable that can be allowed to change is the concentration of the reactant for which the order is being determined. The concentrations of all other reactant species must remain the same. Furthermore, the reaction temperature and the physical state of the reac- tants must be held constant for all trials.

The rate-law expression can be written once the numerical values of the reaction order exponents (n and m) are known.

R = k[A]n[B]m R = k[A]1[B]2

In our example, the reaction order with respect to [A] is 1, and the reaction order with respect to [B] is 2. The overall order of the reaction is the sum of the order for each reactant, which in this example is 1 + 2 = 3.

Once the reaction orders have been experi- mentally determined, the rate constant k can be determined by using data from any of the trials. For example, using the data from Trials 1 and 3 give the same rate constant.

Trial 1 R = k[A][B]2 k = R / [A][B]2 k = (0.04 M/s) / [0.10 M][0.10 M]2

k = (0.040 M/s) / [0.10 M][0.01 M2] k = (0.040 M/s) / [0.001 M3] k = 40 M-2 s-1

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Needed but not supplied: • Ice • Pure water* • Warm water • Timing device or stopwatch

Materials Included in the materials kit:

Hydrochloric acid solution, 0.1 M, HCl

Syringe, 1 mL

Starch solution, 1%

24-Well plate

Potassium iodide solution, 0.05 M, KI

Sodium thiosulfate solution, 0.01 M, Na2S2O3

Needed from the equipment kit:

Test tube rack

Plastic pipets

Graduated cylinder, 10 mL

Plastic cups, 10 oz

Polystyrene test tubes

Thermometer

Plastic medicine cup

Copper(II) sulfate solution, CuSO4

Hydrogen peroxide, 3%, H2O2

Needed from the chemical kit #1:

Needed from the chemical kit #2:

Reorder Information: A replacement kit for Factors Affecting Reaction Rates, item number 580318, can be ordered from Carolina Biological Supply Company. Call 800-334-5551 to order.

*Bottled or purified water should be used because they lack most of the impurities that are found in tap water. Water that has been filtered through a water purifier (e.g., Brita® or PUR®) works well.

FACTORS AFFECTING REACTION RATES

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Preparation 1. Read through the Activity sections to

familiarize yourself with all steps before you begin the experiment.

2. Obtain all required materials. 3. Fill a plastic medicine cup with pure water.

Safety Wear your goggles, gloves, and lab apron at all times while conducting this investigation.

Read all instructions for this laboratory activity before beginning. Follow the instructions closely and observe established laboratory safety practices, including use of appropriate personal protective equipment (PPE).

Copper sulfate is harmful if swal- lowed. Hydrogen peroxide can cause serious eye damage. Potassium iodide may cause an allergic skin reaction. Hydrochloric acid causes skin and eye irritation.

Hydrogen peroxide causes skin irritation.

Copper sulfate is very toxic to aquatic life.

Hydrochloric acid can be toxic if inhaled.

Do not eat, drink, or chew gum while performing this activity. Wash your hands with soap and water before and after performing the activity. Clean up the work area with soap and water after completing the investigation. Keep pets and children away from lab materials and equip- ment.

In this investigation, use a different pipet for each reagent. Wrap a piece of tape around the barrel of the pipet and write the reagent name on the tape. Use a different well in the well plate for each reaction. Do not cross- contaminate the contents of one well to another, and do not allow microdroplets of the solutions to splash into multiple wells during delivery. Use a different row of the plate for each activity. Start at the left well and proceed right. Solutions can be discarded after each activity. Before reusing the plate, rinse all wells several times with purified water and allow them to dry thoroughly.

1 drop for each trial, until it takes between 18–22 seconds for the dark blue color to appear.

13. If the reaction took less than 18–22 seconds, then the amount of sodium thiosulfate must be increased. Repeat steps 1–10 in the next well in the row, but add 9 drops of Na2S2O3 instead of 8 drops. Continue the trials, adding 1 more drop of Na2S2O3 for each trial, until it takes between 18–22 seconds for the dark blue color to appear.

14. Record the number of drops of Na2S2O3 used to achieve the approximately 20-second change in Data Table 1. This is the number of drops that will be used in all subsequent experiments.

ACTIVITY 2 A Concentration

Concentration of KI Perform steps 1–10 described in the calibration activity, but use the number of drops of each reactant that are specified in the following table. Use the number of drops of Na2S2O3 determined in Activity 1 for all trials. In this set of trials, only the amounts of KI and water are changed. The water is added to ensure that the final volume and the concentrations of all reactants are constant.

1. Record the reaction times for these trials in Data Table 2a.

2. Calculate the average reaction time for similar trials (i.e., trials 1 and 2)

3. Calculate the rate for similar trials, 1/Time.

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ACTIVITY 1 A Calibration Tests

1. Orient the 24-well plate so that there are four wells across (horizontal axis) and six wells down (vertical axis).

2. In the top-left well, place 8 drops of KI. 3. Add 2 drops of HCl to the KI solution in the

well. 4. Add 4 drops of starch to the KI solution in the

well. 5. Add 8 drops of Na2S2O3 solution (sodium

thiosulfate) to the KI solution in the well. 6. Mix the solution by gently swirling the 24-well

plate in a small circle on the table. 7. With the syringe, draw up 0.4 mL of H2O2. 8. As the H2O2 is added to the solution in

the well, begin timing.

9. Mix the solution by gentle swirling, and monitor the color.

10. Record the amount of time it takes for the entire solution to turn a dark blue or black color in Data Table 1.

11. If the solution took 18–22 seconds to turn dark blue, then no additional calibration trials are necessary.

12. If the reaction took more than 18–22 seconds, then the amount of sodium thiosulfate must be decreased. Repeat steps 1–10 in the next well in the row, but add 7 drops of Na2S2O3 instead of 8 drops. Continue the trials, reducing the Na2S2O3 by

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ACTIVITY

Add the hydrogen peroxide quickly but without splashing.

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ACTIVITY 2 continued

ACTIVITY

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out the syringe and add the required amount of water to all wells before starting any trials.

1. Record the reaction times for these trials in Data Table 2b.

2. Calculate the average reaction time for similar trials (i.e., trials 1 and 2)

3. Calculate the rate for similar trials, 1/Time.

H2O2 Concentration Effects These trials will follow the same procedure as specified above, but this set of trials will vary the concentration of H2O2. As stated above, the amount of water added to each reaction must be adjusted so that the final volume and the concentrations of all reactants are constant. To avoid contamination, it may be best to rinse

KI trial KI (drops)

HCl (drops)

Starch (drops)

Water (drops)

Na2S2O3 (drops)

H2O2 (mL)

1 8 2 4 0 Activity 1 0.4 2 8 2 4 0 Activity 1 0.4 3 6 2 4 2 Activity 1 0.4 4 6 2 4 2 Activity 1 0.4 5 4 2 4 4 Activity 1 0.4 6 4 2 4 4 Activity 1 0.4 7 2 2 4 6 Activity 1 0.4 8 2 2 4 6 Activity 1 0.4

H2O2 trial

KI (drops)

HCl (drops)

Starch (drops)

Water (mL)

Na2S2O3 (drops)

H2O2 (mL)

1 8 2 4 0.0 Activity 1 0.4 2 8 2 4 0.0 Activity 1 0.4 3 8 2 4 0.1 Activity 1 0.3 4 8 2 4 0.1 Activity 1 0.3 5 8 2 4 0.2 Activity 1 0.2 6 8 2 4 0.2 Activity 1 0.2 7 8 2 4 0.3 Activity 1 0.1 8 8 2 4 0.3 Activity 1 0.1

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Concentration of KI

H2O2 Concentration Effects

7. Calculate the average reaction time for similar trials (i.e., trials 1 and 2)

8. Calculate the rate for similar trials, 1/Time. 9. Rinse well plate three times with bottled

water. Then clean using soap and water and finally rinse again three times with bottled water. Let air dry or blot dry using clean paper towels.

ACTIVITY 3 A Temperature

1. Place one test tube into each of the three plastic cups and allow the tube to lean against the edge. Mark the cups with a horizontal line at the halfway points of the test tubes, and remove the test tubes from the cups.

2. Prepare three water baths as follows: a. For a room-temperature water bath,

add room-temperature water to one cup until the water reaches the mark. Room temperature is approximately 25 °C. If

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Concentration of HCl 1. This trial follows the same procedure as

specified for determining the concentration of KI (see the following table), but it varies the concentration of HCl.

2. To vary the HCl concentration, a series of new HCl concentrations will be prepared from the stock 0.1 M HCl reagent (the 0.1 M HCl stock reagent is the 100% HCl solution). The same volume of HCl will be used in each trial (2 drops), but the HCl concentration contained in this volume will change.

3. In the first well of a row, add 15 drops of 0.1 M HCl and 5 drops of pure water. Swirl to mix. This is the 75% HCl solution.

4. In the second well of the row, add 10 drops of HCl and 10 drops of pure water. Swirl to mix. This is the 50% HCl solution.

5. In the third well of the row, add 5 drops of HCl and 15 drops of pure water. Swirl to mix. This is the 25% HCl solution.

6. Record the reaction times in Data Table 2c.

HCl (drops)

KI (drops)

% HCl (2 drops)

Starch (drops)

Water (drops)

Na2S2O3 (drops)

H2O2 (mL)

1 8 100% 4 0 Activity 1 0.4 2 8 100% 4 0 Activity 1 0.4 3 8 75% 4 0 Activity 1 0.4 4 8 75% 4 0 Activity 1 0.4 5 8 50% 4 0 Activity 1 0.4 6 8 50% 4 0 Activity 1 0.4 7 8 25% 4 0 Activity 1 0.4 8 8 25% 4 0 Activity 1 0.4

Concentration of HCl

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ACTIVITY 3 continued

ACTIVITY

same as those used in the first trial in Activity 2 (Concentration of KI). See the table on top of page 10 for the exact amounts needed.

4. Add 0.4 mL H2O2 to a clean test tube, and place this tube in the appropriate water bath for 5 minutes. This serves to pre-equilibrate the H2O2 solution to the required temperature before adding it to the reaction.

5. After preparing the reaction solutions (step 3 of this activity), gently swirl the solution so the reactants are mixed and the solution reaches the temperature of its water bath.

6. When the reactants have reached the temperature of the water bath, add the pre-equilibrated H2O2 to the reaction test tube and begin timing. The reaction test tube should be kept in the water bath.

7. Swirl the solution and monitor the color. 8. Record the amount of time it takes for the

room-temperature reactants to turn blue- black in Data Table 3.

9. Repeat the mixing and reaction processes with the pairs of tubes in both the hot-water and cold-water baths. Record the time it takes for the reactions to reach a blue-black color in Data Table 3.

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the water temperature varies greatly from this temperature, add a small amount of warm or cool water until it reaches this temperature. Record the temperature in Data Table 3.

b. For a cold-water bath, add ice to a cup until the ice is approximately halfway to the mark and then add water until the level reaches the mark. The cold-water bath should be approximately 5 °C, or approximately 20 °C below that of the room-temperature bath. Immediately before using the cold-water bath, record its temperature in Data Table 3.

c. For a hot-water bath, add hot tap water to a cup until the water level reaches the mark. Ensure that the temperature of this water is approximately 45 °C, or approximately 20 °C above that of the room-temperature bath. Immediately before using the hot-water bath, record its temperature in Data Table 3.

3. In this activity, the KI, HCl, starch, and Na2S2O3 solutions will be added to one test tube and placed in the appropriate water bath for 5 minutes. The amounts of each reactant are the

HCl Trial

KI (drops)

HCl (drops)

Starch (drops)

Water (drops)

Cu2So4 (drops)

Na2S2O3 (drops)

H2O2 (mL)

1 8 2 4 4 0 Activity 1 0.4 2 8 2 4 3 1 Activity 1 0.4 3 8 2 4 2 2 Activity 1 0.4 4 8 2 4 1 3 Activity 1 0.4 5 8 2 4 0 4 Activity 1 0.4

Catalyst

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2. The rate law in terms of the reactants is given by

R = k[I-]m[H2O2]n[H+] p

3. For each reactant, determine the order with respect to that reactant using the following formula:

concentration 1 X = rate 1 concentration 2 rate 2

4. In the equation, x represents the corresponding exponent in the rate law (i.e., n, m, or p). Refer to the background for specifics on this calculation. Record the results Data Table 5.

5. For each reactant, perform the calculation 3 times from highest concentration to lowest concentration, i.e., compare 1 and 2, 2 and 3, and 3 and 4.

6. Determine the average reaction order for each reactant and record them in Data Table 5.

7. Round the average to the nearest whole integer (i.e., 0, 1, 2, 3…) and record these as the final values for m, n, and p.

8. Take the rate law expression for this reaction, R = k[I–]m[H2O2]n[H+]p, and insert the exponents for m, n, and p that you determined in the preceding steps. If any of the experimentally determined exponents is 0 (i.e., no change in rate after doubling in concentration), then mathematically the value will be 1 and the exponent will fall out of the rate equation. Record the rate law expression in Data Table 5.

9. Determine the overall order of this reaction by finding the sum of all the exponents. Record this in Data Table 5.

ACTIVITY 4 A Catalyst

1. Add one drop of 0.5 M copper(II) sulfate solution to one of the wells.

2. Add about 3 mL of pure water to the well to dilute the copper sulfate. Swirl the solution to mix.

3. Repeat the reaction with the following reactant quantities. Use the dilute copper sulfate solution that was just prepared.

4.

Record the time in Data Table 4.

Disposal and Cleanup 1. Dispose of reaction solutions in the sink, and

then rinse the sink well. 2. Clean and dry all equipment and return it to

the equipment kit. 3. Clean the work space.

ACTIVITY 5 A Determining the Rate Law

1. The rate law is determined by observing the effect of different concentrations of I-, H2O2, and HCl on the rate of the reaction (Activity 2). The exponents (reaction order) of each reactant are determined by evaluating the change in reaction rate when only that reactant concentration is changed. The overall chemical reaction is written as follows:

3I‾ (aq) + H2O2 (aq) + 2H+ (aq) → I3‾ (aq) + 2HO2 (l)

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DATA TABLES

ACTIVITY

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Data Table 1. Calibration

Data Table 2a. Effects of KI Concentration

Data Table 2b. Effects of H2O2 Concentration

Trial Na2S2O3 (Drops) Reaction Time (sec)

1 2 3

How many drops of Na2S2O3 will be used in the remaining experiments?

____________________________

Trial KI (drops)

HCl (drops)

Starch (drops)

H2O (drops)

Na2S2O3 (drops)

H2O2 (mL)

Time1 (sec)

Time2 (sec)

Average Time

(Tavg) (sec)

Rate (1/Tavg)

1 8 2 4 0 0.4 2 8 2 4 0 0.4 3 6 2 4 2 0.4 4 6 2 4 2 0.4 5 4 2 4 4 0.4 6 4 2 4 4 0.4 7 2 2 4 6 0.4 8 2 2 4 6 0.4

Trial KI (drops)

HCl (drops)

Starch (drops)

H2O (mL)

Na2S2O3 (drops)

H2O2 (mL)

Time1 (sec)

Time2 (sec)

Average Time

(Tavg) (sec)

Rate (1/Tavg)

1 8 2 4 0 0.4 2 8 2 4 0 0.4 3 8 2 4 0.1 0.3 4 8 2 4 0.1 0.3 5 8 2 4 0.2 0.2 6 8 2 4 0.2 0.2 7 8 2 4 0.3 0.1 8 8 2 4 0.3 0.1

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Data Table 2c. Effects of H+ Concentration

Data Table 3. Temperature

Data Table 4. Catalyst

Trial KI (drops)

HCl (%) Starch (drops)

H2O (drops)

Na2S2O3 (drops)

H2O2 (mL)

Time1 (sec)

Time2 (sec)

Average Time

(Tavg) (sec)

Rate (1/Tavg)

1 8 100% 4 0 0.4 2 8 100% 4 0 0.4 3 8 75% 4 0 0.4 4 8 75% 4 0 0.4 5 8 50% 4 0 0.4 6 8 50% 4 0 0.4 7 8 25% 4 0 0.4 8 8 25% 4 0 0.4

Water Bath Trial Temperature of the Water Bath (°C)

Reaction Time (sec)

Cold water Room-temp. water Hot water

Trial Water (drops)

CuSO4 (drops)

Reaction Time (sec)

1 4 0 2 3 1 3 2 2 4 1 3 5 0 4

ACTIVITY

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DATA TABLES Data Table 5. Orders of Reactants in the Rate Law

Reactant Concentration Average Rate

Calculated Reaction Order (X)

Average Reaction

Order

Reaction Order

(integer)

I- 8 drops m =

I- 6 drops

I- 4 drops

I- 2 drops

H2O2 0.4 mL n =

H2O2 0.3 mL

H2O2 0.2 mL

H2O2 0.1 mL

HCl 100% p =

HCl 75%

HCl 50%

HCl 25%

Rate Law

Overall Order of Reaction

concentration 1 X = rate 1 concentration 2 rate 2( ) e.g.

R = k[I-]m[H2O2]n[H+] p

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NOTES

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NOTES

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CHEMISTRY Factors Affecting Reaction Rates

Investigation Manual

CB780481703

Carolina Biological Supply Company www.carolina.com • 800.334.5551 ©2017 Carolina Biological Supply Company

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