Chemistry Experiments for Nova5000

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Experiments in Chemistry for the Nova5000

First Edition

First Print

Printed in June 2007

Table of Contents

Experiments by Sensors ............................................................................................. 1

Preface........................................................................................................................ 3

1. Acid - Base Titration: Reaction of NaOH with HCl .................................................. 5

2. Exothermic Reactions - Dissolution of NaOH in Water ......................................... 11

3. Reduction and Oxidation (Redox) Reactions – Copper Chloride with

Aluminum .................................................................................................................. 17

4. Endothermic Reactions - Dissolution of Ammonium Nitrate in Water ................... 23

5. Endothermic Reactions - Mixing Crystals of Barium Hydroxide and

Ammonium Isothiocyanate ........................................................................................ 27

6. Endothermic Reactions – Reaction of Citric Acid Solution with Baking Soda....... 33

7. Chemical Catalysis: Decomposition of H2O2 in the Presence of MnO2 ................. 39

8. Effects of Variations in Air Temperature on Air Pressure - The Combined

Gas Law .................................................................................................................... 47

9. Effects of Variations in Air Volume on Air Pressure – Boyle’s Law....................... 53

10. Hess’ Law: Additivity of Heat of Reactions.......................................................... 59

11. Heat of Combustion............................................................................................. 65

12. Saltwater Conductivity......................................................................................... 71

13. The Lambert-Beer Law........................................................................................ 75

14. The Ideal Gas Law .............................................................................................. 81

15. Exploring a Candle Flame................................................................................... 87

16. Freezing and Melting of Water ............................................................................ 91

17. Another Look at Freezing Temperature .............................................................. 95

18. Energy Contents of Foods................................................................................... 99

19. Chemical Equilibrium: Finding a Constant, Kc .................................................. 105

20. Energy Content of Fuels.................................................................................... 111

Experiments by Sensors 1

Experiments by Sensors

Colorimeter

13. The Lambert-Beer Law ........................................................................................75

19. Chemical Equilibrium: Finding a Constant, Kc ...................................................105

Conductivity

12. Saltwater Conductivity..........................................................................................71

pH

1. Acid - Base Titration: Reaction of NaOH with HCl ...................................................5

2. Exothermic Reactions - Dissolution of NaOH in Water ..........................................11

Pressure (150 –1150 mbar)

7. Chemical Catalysis: Decomposition of H2O2 in the Presence of MnO2................39

8. Effects of Variations in Air Temperature on Air Pressure - The Combined Gas

Law ............................................................................................................................47

9. Effects of Variations in Air Volume on Air Pressure – Boyle’s Law........................53

14. The Ideal Gas Law...............................................................................................81

16. Freezing and Melting of Water .............................................................................91

17. Another Look at Freezing Temperature ...............................................................95

18. Energy Contents of Foods ...................................................................................99

20. Energy Content of Fuels ....................................................................................111

2 Experiments by Sensors

Temperature (-25 ºC to 110 ºC)

1. Acid - Base Titration: Reaction of NaOH with HCl ...................................................5

2. Exothermic Reactions - Dissolution of NaOH in Water ..........................................11

3. Reduction and Oxidation (Redox) Reactions – Copper Chloride with Aluminum ..17

4. Endothermic Reactions - Dissolution of Ammonium Nitrate in Water ....................23

5. Endothermic Reactions - Mixing Crystals of Barium Hydroxide and Ammonium

Isothiocyanate ............................................................................................................27

6. Endothermic Reactions – Reaction of Citric Acid Solution with Baking Soda........33

8. Effects of Variations in Air Temperature on Air Pressure - The Combined Gas

Law ............................................................................................................................47

10. Hess’ Law: Additivity of Heat of Reactions...........................................................59

11. Heat of Combustion .............................................................................................65

Temperature (0 ºC to 1250 ºC)

15. Exploring a Candle Flame....................................................................................87

Preface 3

Preface

This book contains 20 student experiments in Chemistry to be conducted using the

Nova5000 and Fourier sensors.

For your convenience we have added an index in which the experiments are sorted

according to sensors (see page 1).

Experiment Layout Each experiment includes the following parts:

• Introduction

A brief description of the concept

and theory

• Equipment

The equipment needed for the

experiment

• Equipment Setup Procedure

Illustrated guide to assembling the

experiment

• Data Logger Setup

Recommended setup

• Experimental Procedure

Step-by-step guide to executing

the experiment

• Data Analysis

• Questions

• Further Suggestions

Safety Precautions • Follow standard safety procedures for laboratory activities in a science classroom.

• Proper safety precautions must be taken to protect teachers and students during the

experiments described in this book.

• It is not possible to include every safety precaution or warning!

• Fourier assumes no responsibility or liability for use of the equipment, materials, or

descriptions in this book.

4 Preface

Acid - Base Titration: Reaction of NaOH with HCl 5

1. Acid - Base Titration: Reaction of NaOH with HCl

Figure 1

Introduction In aqueous solutions, addition of bases to water leads to an increase in the pH of the

solution, while the addition of acids leads to a decrease in the pH. The changes in the

pH can be followed using either specific dyes, called indicators, or a pH electrode. Acids

and bases neutralize, or reverse, the action of one another. By adding a known amount

6 Acid - Base Titration: Reaction of NaOH with HCl

of acid to a basic solution, until it completely reacts with it, the amount of the base can

be determined. This procedure is called: Acid - Base Titration. During neutralization,

acids and bases react with each other to produce ionic substances, called salts.

In this experiment, changes in pH and temperature, occurring while an acid (hydrochloric

acid) is added to a base (sodium hydroxide) solution, are followed using a pH electrode

and a Temperature sensor.

Equipment • Nova5000

• A pH electrode

• A Temperature sensor (-25 °C to 110 °C)

• A polystyrene coffee cup

• 50 ml buret pipette

• A glass funnel

• 2 g NaOH

• 100 ml NHCl solution

• Magnetic stirrer

Equipment Setup Procedure 1. Launch MultiLab.

2. Connect the pH electrode to Input 1 (I/O-1) of the Nova5000.

3. Connect the Temperature sensor to Input 2 (I/O-2) of the Nova5000.

4. Assemble the equipment as illustrated in Figure 1.

5. Click Setup on the main toolbar and program the data logger according to the

setup specified below.


Data Logger Setup

Sensors:

Input 1: pH

Input 2: Temperature (-25 °C to 110 °C)

Rate:

Every second

Samples:

2000 samples

Experimental Procedure 1. Prepare a polystyrene cover for the polystyrene coffee cup. The cover should be flat

and larger than the circumference of the coffee cup (see Figure 1).

2. Pierce three holes in the cover: One for the pH electrode, one for the Temperature

sensor and one for the glass funnel.

3. Add 50 ml 0.5 N NaOH solution to the coffee cup.

4. Place the coffee cup on a magnetic stirrer.

5. Place the cover on the cup.

6. Click Run on the upper toolbar to begin recording data.

7. Wait until readings from the sensors are stable.

8. Start stirring the NaOH solution in the coffee cup.

9. Add, drop-by-drop, 1 N HCl solution from the buret to the cup, through the glass

funnel.

10. Follow changes in pH and temperature registered on the monitor.

11. As the pH just starts to change, stop the flow of HCl. Find the volume of HCl added

until that point.

12. Renew the flow of HCl. Follow pH changes very carefully.

13. Stop immediately the flow of HCl as the pH level stabilizes.


14. Save your data by clicking Save on the upper toolbar.

Data Analysis

1. Use the First cursor and Second cursor to display the change in the pH

obtained during the neutralization process.

What was the initial pH value? The final value? The difference between the two

values?

The Cursor: You can display up to two cursors on the graph simultaneously.

Use the first cursor to display individual data recording values, to select a curve or to

reveal the hidden Y-axis.

Use two cursors to display the difference between two coordinate values or to select a

range of data points.

To display the First cursor: Double click on an individual data point or click First cursor

on the graph toolbar. You can drag the cursor with the mouse onto any other point

on the plot, or onto a different plot. For finer cursor movements click Forward and

Backward cursor.

The coordinate values of the selected point will appear in the information bar at the

bottom of the graph window.

To display the Second cursor: Double click again anywhere on the graph area or click

Second cursor .

The information bar will now display the difference between the two coordinate values.

To remove the cursors: Click to remove the Second cursor and click to remove

the First cursor.

2. What was the volume of HCl added until the pH started to change? Compare it with

the volume of HCl added until a complete neutralization of NaOH was achieved.

3. Use the graph cursors to find the time interval between the point of start of change in

pH and the point of pH start of stabilization.


4. Use the graph cursors to display the change in temperature obtained in the process

and the time needed to reach the final temperature value.

5. Calculate the heat of reaction: C- water capacity, ∆T - temperature change.

TmcQ ∆=

Note: Water specific heat capacity at 25 °C is 4.18 (J/g*°C).

An example of the graph obtained in this experiment is shown below:

Figure 2

Questions 1. Did you observe a fast pH change? Explain the difference in the short time interval

needed for the completion of the drastic changes of pH and that of the whole

neutralization process.

2. Is the neutralization reaction an exothermic reaction? Base your conclusions on the

experiment you performed.

3. Try to assume the results of acid - base reaction at different concentrations of NaOH

added to the coffee cup. What will be the pH change in each case? What will be the

extent of change in the temperature?

4. What will be the effect of reacting other acids (such as, for example, acetic acid) with

NaOH?

pH

Temp


Further Suggestions 1. React different concentrations of NaOH with a constant concentration of HCl.

2. Calculate unknown concentrations of the titrated NaOH (or HCl). This can be

achieved by setting the flow of acid (or base) from the Buret pipette at a constant

rate (from the time interval on X axis multiplied by the flow rate, the volume added,

can be calculated).

3. Examine the effect of increasing the water and/or the surrounding temperature, on

the reaction.

4. Perform Acid - Base reaction with different types of acids and/or bases: A weak acid

with a strong base and vice versa.

Exothermic Reactions - Dissolution of NaOH in Water 11

2. Exothermic Reactions - Dissolution of NaOH in Water

Figure 1

Introduction Nearly all chemical reactions involve either the release or the absorption of heat. The

reactions are classified as exothermic or endothermic. An exothermic process is a

chemical reaction in which heat is evolved. When we perform an exothermic reaction in

a flask, it initially warms. Later, heat from the flask flows to the surrounding air until a

12 Exothermic Reactions - Dissolution of NaOH in Water

temperature balance is established. A calorimeter is the device used to measure the

heat absorbed or evolved during a chemical reaction.

The heat of reaction can be calculated from the following equation:

q = C∆t

Where q = quantity of heat evolved/absorbed, C = heat capacity, and ∆t = temperature

change.

In this experiment we follow temperature changes occurring during dissolution of sodium

hydroxide in water. A polystyrene coffee cup will serve as a calorimeter.


• A pH electrode

• A Temperature sensor (-25 °C to 110 °C)

• A polystyrene coffee cup

• 10 g NaOH



2. Connect the pH sensor to Input 1 (I/O-1) of the Nova5000.






Data Logger Setup

Sensors:

Input 1: pH


Rate:

Every second

Samples:

5000 samples

Experimental Procedure 1. Prepare a polystyrene cover to the polystyrene coffee cup. The cover should be flat

and larger than the circumference of the coffee cup. Pierce two holes in the cover;

one for the pH electrode, the other for the Temperature sensor.

2. Pour 100 ml of tap water into the coffee cup.


4. Place the cover on the cup, but leave a narrow opening so that NaOH can be added.



7. Start stirring the water in the coffee cup.

8. Add 2 g crystalline NaOH to the cup and immediately adapt the cover so that the cup

will be well covered.

9. Follow changes in pH and temperature registered on the monitor, until no further

changes in either pH or temperature are observed.

10. Click Stop on the upper toolbar to stop collecting data.

11. Repeat the dissolution reaction, this time in a transparent beaker (this part of the

experiment can be performed in parallel to the first part, using two pH electrodes and

two Temperature sensors).


12. Add 100 ml water to the beaker and insert a pH electrode and Temperature sensor

into it.

13. Start stirring the water.


15. Add 2 g of solid NaOH to the water.

16. Follow the dissolution of the NaOH in the water until it is completed.



Data Analysis

1. Use the First cursor and Second cursor to display the change in the pH

obtained during the dissolution process.

What was the initial pH value? The final value? The difference between the two

values?









Backward cursor.




Second cursor .




the First cursor.

2. Use the graph cursors to find:

a. The time needed to reach the final value of pH.

b. The change in temperature obtained in the process and the time needed to reach

the final temperature value.

c. The time interval needed to complete the dissolution of NaOH.

3. Calculate the heat of reaction: C- water capacity, ∆t - temperature change.



Figure 2

Questions 1. Did you observe a fast pH change? Compare the time span for completion of the pH

changes with that of the temperature.

pH

Temp


2. Explain the difference in the time span needed for the completion of pH changes and

that of the temperature.

3. Is the dissolution of NaOH an exothermic reaction? Is it a vigorous reaction? Base

your conclusions on the experiment you performed.

4. Try to guess what the results of the dissolution would be according to different

amounts of NaOH in water. What will the pH change be in each case? What will be

the extent of change in the temperature?

Further Suggestions 1. Dissolve different amounts of NaOH in water. Follow changes in pH and temperature

in each case. Calculate heat of the reaction in each case.

2. Examine the effect of increasing the water and/or the surroundings temperature, on

the dissolution of NaOH.

3. Examine the effect of the water pH on the dissolution of NaOH. Follow the heat of

reaction in a buffer solution. Alternatively, dissolve in the water KOH or NH4OH, prior

to the dissolution of NaOH.

4. Perform an additional exothermic reaction. Dissolve anhydrous CuSO4 (white

crystals) in water. Dissolution of Copper Sulfate in water forms a blue hydrated

copper ion.

Reduction & Oxidation (Redox) Reactions – Copper Chloride with Aluminum 17

3. Reduction and Oxidation (Redox) Reactions – Copper Chloride with Aluminum

Figure 1

Introduction Redox reactions primarily involve the transfer of electrons between two chemical

species. The compound that loses an electron is said to be oxidized, the one that gains

an electron is said to be reduced. There are also specific terms that describe the

18 Reduction & Oxidation (Redox) Reactions – Copper Chloride with Aluminum

specific chemical species. A compound that is oxidized is referred to as a reducing

agent, while a compound that is reduced is referred to as the oxidizing agent.

The oxidation number is a number identical with the valence but with a sign, expressing

the nature of the charge of the species in question when formed from the neutral atom.

Thus, the oxidation number of chlorine in hydrochloric acid is -1, while it is +1 in

hypochlorous acid. Similarly we can say that the oxidation number of chlorine in chloric

acid (HCIO3) is +5, and in perchloric acid (HCIO4) +7.

In this experiment we will be following the temperature changes that occur during redox

reaction:

3Cu2+(aq)+6Cl2-

2 (aq)+2Al0(s) 3Cu0(s)+2Al3+

(aq)+6Cl2-(aq)

Aluminum, Al, was oxidized to Al3+, Cu2+ was reduced back to Cu.


q = C∆t


change.


• Temperature sensor (-25 °C to 110 °C)

• Polystyrene coffee cup

• 5 g CuCl2


• Aluminum paper

• Safety goggles







Data Logger Setup

Sensors:


Rate:

Every second

Samples:

200 samples

Experimental Procedure 1. Obtain and wear safety goggles.

2. Prepare a polystyrene cover to the polystyrene coffee cup. The cover should be flat

and larger than the circumference of the coffee cup. Pierce a hole in the cover for the

Temperature sensor.



5. Place the cover on the cup but leave a narrow opening so that CuCl2 can be added.

6. Click Run on the upper toolbar to begin recording data. Wait until readings from

the sensors are stable.


8. Add 5 g crystalline CuCl2 to the cup and immediately replace the cover so the cup


9. Add the Aluminum paper to the cup.

10. Follow changes in temperature registered on the monitor, until no further changes in

temperature are observed.




Data Analysis

1. Use the First cursor and the Second cursor to display the change in the

process and the time needed to reach the final temperature value.









Backward cursor.




Second cursor .



the First cursor.

2. Calculate the heat of reaction.



An example of the graph Temperature vs. Time obtained in this experiment is shown

below:

Figure 2

Questions 1. What changes occurred in the color of the Aluminum paper?

2. Write the reaction that took place and write the redox equation. Which atom has

reduced, and which was oxidized?

Further Suggestions 1. Drop a piece of iron into a solution of copper chloride. Write the reaction that occurs,

and the redox equation.

2. Drop a piece of Zinc metal into dilute hydrochloric acid. Hydrogen gas, H2, is given

off. Write the general equation and redox equation. Which atom was reduced?

Which atom was oxidized?


Endothermic Reactions - Dissolution of Ammonium Nitrate in Water 23

4. Endothermic Reactions - Dissolution of Ammonium Nitrate in Water

Figure 1

Introduction An endothermic process is a chemical reaction in which heat is absorbed. When we

perform an endothermic reaction in a flask, it initially cools. Later, heat from the

surroundings flows to the flask until the temperature balance is established.

In this experiment we follow temperature changes occurring during the dissolution of

crystalline ammonium nitrate in water.

24 Endothermic Reactions - Dissolution of Ammonium Nitrate in Water

The heat of the reaction can be calculated from the following equation:

q = C∆t


change.



• 5 g NH4NO3








Data Logger Setup

Sensors:


Rate:

Every second

Samples:

500 samples

Endothermic Reactions - Dissolution of Ammonium Nitrate in Water 25

Experimental Procedure 1. Prepare a polystyrene cover to the polystyrene coffee cup. The cover should be flat

and larger than the circumference of the coffee cup. Pierce a hole in the cover, and

insert the Temperature sensor.



4. Place the cover on the cup but leave a narrow opening so that NH4NO3 can be

added.




8. Add 5 g crystalline NH4NO3 to the cup and immediately replace the cover to ensure

that the cup is well covered.

9. Follow changes in temperature registered on the monitor until no further changes in




Data Analysis


obtained temperature in the process and the time needed to reach the final

temperature value.

2. Calculate the heat of the reaction.

Note: Water specific heat capacity at 25 °C is 4.18 (J/g* °C).

26 Endothermic Reactions - Dissolution of Ammonium Nitrate in Water

An example of the graph Temperature vs. Time obtained in this experiment is shown

below:

Figure 2

Questions 1. What kind of chemical reaction is the dissolution of ammonium nitrate in water? Base

your conclusions on the experiment you performed.

2. Try to assume the results of the dissolution of different amounts of NH4NO3 in water.

What will be the extent of change in the temperature?

3. What will be the effect of warming the water prior to the dissolution of NH4NO3 in it?

What will be the effect of cooling the water?

Further Suggestions 1. Dissolve different amounts of NH4NO3 in water. Follow changes in temperature in

each case. Calculate heat of the reaction in each case.

2. Examine the effect of increasing the water and/or the surrounding temperature, on

the dissolution of NH4NO3.

Endothermic Reactions - Mixing Crystals of Barium Hydroxide and Ammonium

Isothiocyanate 27

5. Endothermic Reactions - Mixing Crystals of Barium Hydroxide and Ammonium Isothiocyanate

Figure 1A

Introduction When the two crystalline substances Ba(OH)2 × 8HO and NH4SCN are thoroughly mixed

in a flask, a heat absorbing reaction is developed:

28 Endothermic Reactions - Mixing Crystals of Barium Hydroxide and

Ammonium Isothiocyanate

O)10H2NHBa(SCN) 2NH4SCNOH8Ba(OH) 23222 ++→+×

The gaseous substance Ammonia (NH3), formed in this reaction, can be easily detected

due to its pungent smell.

If the flask, which feels very cold to the touch, is set on a board covered with a thin layer

of water, the flask and the board freeze solidly together.

In this experiment we follow temperature changes occurring during the mixing of

crystalline barium hydroxide octahydrate and ammonium Isothiocyanate. We experiment

the freezing together of the reaction flask and the board.


q = C∆t


change.



• 2 g of Ba(OH)2 × 8H2O

• 4 g of NH4SCN

• Wooden or a plastic board, approximately 5 cm x 5 cm

• 10 ml glass flask

• 10 cm glass rod



3. Assemble the equipment as illustrated in Figure 1A.




Isothiocyanate 29

Data Logger Setup

Sensors:


Rate:

Every second

Samples:

200 samples

Experimental Procedure 1. Pour some water on the board until it is partly covered with a thin layer.

2. Weigh 2 g of Ba(OH)2 × 8H2O in a 10 ml glass flask.

3. Insert the Temperature sensor into the barium hydroxide crystals in the flask.

4. Weigh 4 g of ammonium Isothiocyanate.


6. Wait until readings from the Temperature sensor are stable.

7. Add the weighed NH4SCN to the flask containing Ba(OH)2 × 8H2O.

8. Place the flask on the board covered with the thin layer of water, as demonstrated in

Figure 1B.



Figure 1B

9. Mix thoroughly the substances in the flask with the glass rod.

10. Follow changes in temperature in the flask, registered on the monitor, until no further

changes in temperature are observed.

11. When the temperature is stabilized, click Stop on the upper toolbar to stop

collecting data.


13. Try to remove the flask from the board.


Isothiocyanate 31

Data Analysis

1. Use the First cursor and the Second cursor to display the change in

temperature obtained in the process and the time needed to reach the final

temperature value.









Backward cursor.




Second cursor .



the First cursor.

2. Calculate the heat of reaction.

Note: Water specific heat capacity at 25 °C is 4.18 (J/g×°C)

3. Summarize the results obtained with the board.



An example of the Temperature vs. Time graph obtained in this experiment is shown in

the graph below.

Figure 2

Questions 1. What changes in the temperature in the flask did you observe? Explain your results.

2. What kind of chemical reaction occurred in the flask?

3. Explain the results you obtained with the board.

Further Suggestions 1. Change the relative amounts of Ba(OH)2 × 8H2O and NH4SCN and follow the

temperature change in each case.

2. Perform an additional endothermic reaction: Dissolution of KNO3 in water (25 g in 50

ml water).

3. Follow the rate of ammonium release in the reaction, using a Pressure sensor. In this

way the reaction rate can be measured.

Endothermic Reactions – Reaction of Citric Acid Solution with Baking Soda 33

6. Endothermic Reactions – Reaction of Citric Acid Solution with Baking Soda

Figure 1

Introduction An endothermic process is a chemical reaction in which heat is absorbed. When we

perform an endothermic reaction in a flask, it initially cools. Later, heat from the

surroundings flows to the flask until temperature balance is established.

In this experiment we follow temperature changes occurring during the reaction between

citric acid solution and baking soda.

H3C6H5O7(aq)+3NaHCO3(s) 3CO2(g)+3H2O(l)+NaC6H5O7(aq)

34 Endothermic Reactions – Reaction of Citric Acid Solution with Baking Soda


q = C∆t


change.


• Safety goggles


• 25 ml citric acid H3C6H5O7 solution

• 15 g baking soda NaHCO3









Data Logger Setup

Sensors:


Rate:

Every second

Samples:

500 samples

Experimental Procedure 1. Obtain and wear safety goggles.



insert the Temperature sensor.

3. Pour 25 ml citric acid to the coffee cup.


5. Place the cover on the cup, but leave a narrow opening so that the baking soda

(NaHCO3) can be added.

6. Click Run on the upper toolbar to begin recording data. Wait until readings from

the sensors are stable.


8. After 20 seconds add the baking soda (NaHCO3) to the cup and immediately replace

the cover to ensure that the cup is well covered.






Data Analysis


temperature obtained in the process and the time needed to reach the final

temperature value.









Backward cursor.




Second cursor .



the First cursor.

2. Calculate the heat of the reaction.


An example of the Temperature vs. Time graph obtained in this experiment is shown below:

Figure 2

Questions 1. What kind of chemical reaction is obtained by mixing citric acid solution and baking

soda? Base your conclusions on the experiment you performed.

2. Try to assume the results of mixing different amounts of baking soda in a citric acid

solution. What will be the extent of change in the temperature?

Further Suggestions 1. Change the relative amounts of citric acid and baking soda. Follow changes in

temperature in each case. Calculate the heat of reaction in each case.

2. Follow the rate of CO2 release in the reaction, using the CO2 sensor.


Chemical Catalysis: Decomposition of H2O2 in Presence of MnO2 39

7. Chemical Catalysis: Decomposition of H2O2 in the Presence of MnO2

Figure 1

Introduction Catalysis is the increase in rate of a reaction as the result of addition of a substance

referred to as a catalyst. A catalyst is a substance which takes part in a reaction but is

not consumed by this reaction. The catalyst enters the reaction in one step and is

regenerated in another step.

40 Chemical Catalysis: Decomposition of H2O2 in Presence of MnO2

A pure solution of H2O2 is stable. But when a catalyst, such as MnO2, platinum metal or

Fe+2 ions are added, it rapidly decomposes to water and molecular oxygen.

In this experiment we follow the release of molecular oxygen resulting from the

decomposition of H2O2 in the presence of MnO2, using Pressure sensors.


• Two 10 ml glass bottles

• Two rubber corks, one for each of the bottles

• Two ml plastic syringe

• Two 20-gauge syringe needles

• Three short latex tubes

• 3% H2O2 solution

• Crystal graph MnO2

• Two Pressure sensors (150 – 1150 mbar)

• Two three-way valves


2. Connect the Pressure sensors to Input 1 (I/O-1) and Input 2 (I/O-2) of the Nova5000.

3. Assemble the equipment as illustrated in Figure 1 and Figure 2.

MnO2

2H2O2 2H2O + O2


A syringe needle (20-gauge) is inserted through the cork, until its tip projects slightly out

of the cork (see Figure 2).

Figure 2

At the other end of the syringe, projecting out of the upper side of the cork, a three-way

valve (the one used in infusions) is attached via a very short latex tube. Its length must

be just enough to hold together the valve and the syringe. A Pressure sensor is

connected to the valve, through another short latex tube.

Turn the valve until its opening is directed vertically. In this position, air can flow through

the valve. In order to stop airflow, turn the valve until its opening reaches a horizontal

position.

Into one of the corks, insert an additional needle. A syringe filled with 6% H2O2 solution

will later be attached to this needle.




Data Logger Setup

Sensors:

Input 1: Pressure (150 – 1150 mbar)


Rate:

Every second

Samples:

500 samples

Experimental Procedure 1. Mark the bottles with labels 1 and 2.

2. Fill the plastic syringe with 2 ml of 3% H2O2 solution.

3. Add to bottle 1, 8 ml water and 2 ml 3% H2O2 solution.

4. Add to bottle 2, 8 ml water and a few crystals of MnO2. Gently mix the solution.

5. Close the bottles tightly with the rubber corks.

6. Attach to bottle 2, the syringe filled with H2O2 solution, through the additional needle

inserted through the cork.


8. Follow the pressure level registered on the computer monitor.

9. Turn the valves attached to the corks of the two bottles, until an atmospheric

pressure is obtained in both of them. (0 level is an atmospheric pressure, about 1000

mbar).

10. Inject the H2O2 solution into bottle 2 and immediately turn the valves of the two

bottles, to stop airflow through them.

11. Follow changes in pressure registered on the computer monitor during the

experiment.




Data Analysis

Use the First cursor and the Second cursor to display the change in pressure

in each of the bottles. What was the initial pressure value? What is the final value?

What is the difference between the two values?









Backward cursor.




Second cursor .



the First cursor.

1. Calculate the reaction rate of H2O2 decomposition. Create a differential graph:

subtract the Pressure graph obtained in the control system from that of the

experimental system:

a. Click Functions on the main toolbar.

b. In the Functions drop-down menu select Subtract.


c. In the G1 drop-down menu, select Pressure I/O-1. In the G2 drop-down menu

select Pressure I/O-2.

d. In the Name text box enter a name (e.g. Difference).

e. Click OK.

2. Apply a linear fit to the difference graph:

a. Use the cursors to select the desired range.

b. Click Linear fit on the main toolbar. The fit equation will be displayed in the

information bar at the bottom of the graph window.

c. The slope of the fit line is the net reaction rate.


Figure 3

Questions 1. How is the pressure developed in the experiment related to decomposition of H2O2?

2. Compare the changes in pressure in the two bottles. Did you observe a change in

bottle 1? In bottle 2? Explain the differences.

3. Which of the bottles serves as a control? Explain.


4. Why is a control system needed in the experiment?

5. What is the effect of adding MnO2 crystals to the experimental bottles?

6. What can be the effect of adding increasing amounts of MnO2 on the reaction rate?

7. What can be the effect of temperature rise in the bottles during the experiment, on

the measurement of H2O2 decomposition rate?

Further Suggestions 1. Add increasing amounts of MnO2 to the reaction mixture and follow the reaction

developed in each case.

2. Calculate the reaction rate obtained in each experiment.

3. Compare the effect of different types of chemical catalysts: HBr, HI, Fe+2 ions,

platinum metal.

4. Change the concentration of H2O2 added to the reaction mixture. Compare the effect

of reactant concentrations on the reaction rate, with that of the catalyst.

5. Follow temperature changes occurring during the reaction. Evaluate the effect of

temperature on the decomposition rate of H2O2.


Effects of Variations in Air Temperature on Air Pressure - The Combined Gas Law 47

8. Effects of Variations in Air Temperature on Air Pressure - The Combined Gas Law

Figure 1

Introduction The volume of gases is affected by temperature. As stated in the Charles-Gay Lussac

Law, a sample of gas at a fixed pressure increases in volume linearly with temperature:

V ∝ T or V/T = constant

48 Effects of Variations in Air Temperature on Air Pressure - The Combined Gas Law

Charles Law combined with Boyle’s Law can be expressed in one statement - the

Combined Gas Law. This law states that the volume occupied by a given amount of gas

is proportional to the absolute temperature divided by the pressure:

PV/T = constant

In this experiment we investigate the relationships existing between pressure and

temperature in their effect on gas behavior, by measuring the effect of warming a

constant volume of air trapped in a sealed flask on its pressure.



• Rubber cork



• Three-way valve


• Pressure sensor (150 – 1150 mbar)

• Stand



2. Connect the Pressure sensor to Input 1 (I/O-1) of the Nova5000.


4. Assemble the equipment as illustrated in Figure 2 below.

A syringe needle (20-gauge) is inserted through the cork, until its tip projects slightly

out of the cork (Figure 2).


Figure 2

At the other end of the needle, projecting out of the upper side of the cork, a three-

way valve (the one used in infusions) is attached via a very short latex tube. A

Pressure sensor is attached to the latex tube on the other end.

Turn the valve until its opening is directed horizontally. In this position, air can flow

through the valve from the flask to the surroundings. In order to stop airflow, turn the

valve until its opening reaches a vertical position. In this position air can flow from the

flask to the Pressure sensor.



Data Logger Setup

Sensors:



Rate:

Every second

Samples:

500 samples


Experimental Procedure 1. Assemble the equipment as illustrated in Figure 1 above.

2. Pierce a hole in the cork, fitted in diameter to the tip of the Temperature sensor.

Alternatively, make a narrow slot on the side of the cork.

3. Fill the glass flask with water. Leave a small volume of air in the flask, to prevent

water penetration into the needles.

4. Adapt the Temperature sensor to the hole or to the slot prepared in the cork.

5. Close the cork and start stirring the water in the flask.


7. Follow changes in pressure registered on the screen. Turn the valve, if necessary, to

maintain pressure levels in the flask at atmospheric level before the experiment is

started. Then, turn the valve to its initial position.

8. Start warming the flask. Turn the warming button of the magnetic stirrer to a middle

position. Follow changes in pressure for about 5 minutes.



Data Analysis 1. Compare the pattern of changes in the pressure with that of temperature: Did you

find any similarity between the two of them?

An example of the original and displayed graphs obtained in this experiment, are

shown below:


Figure 3

2. Find the correlation between the changes in temperature and pressure, display a

graph of the Pressure vs. Temperature:

a. Click Format graph on the graph toolbar.

b. Select Temperature I/O-1 in the X-axis list, and then click OK.

An example of the Pressure vs. Temperature obtained in this experiment is shown

below:

Figure 3

Pressure

Temperature


Questions 1. Changes in temperature and in pressure displayed a non-linear curve. Explain this

behavior.

2. What would be the shape of the curves if the flask was warmed in a bath?

3. How is the pressure developed in the experiment related to the changes in

temperature following warming of the flask? In your answer, relate to the Pressure

vs. Temperature graph you performed.

4. What would be the effect of cooling the flask on the pressure in it?

5. Suppose that the water volume in the flask is reduced. What will be the effect of

warming the flask, on the pressure, compared to the results in the present

experiment?

Further Suggestions 1. Warm the flask for a while. Then, stop warming the flask. As the temperature

stabilizes, start cooling the flask. Follow the changes in pressure measured in the

flask

2. Perform the experiment with different water volume in the flask. Compare the effect

of warming or cooling on the pressure in each instance.

Effects of Variations in Air Volume on Air Pressure – Boyle’s Law 53

9. Effects of Variations in Air Volume on Air Pressure – Boyle’s Law

Figure 1

Introduction Gases are distinguished from liquids and solids by their ability to be compressed into

much smaller volumes. According to Boyle’s Law, the volume of a sample of gas at a

given temperature varies inversely with the applied pressure:

P∝1/V or PV = constant

54 Effects of Variations in Air Volume on Air Pressure – Boyle’s Law

In this experiment we investigate the relationships existing between pressure and

volume in their effect on gas behavior, by measuring the effects of changes in air volume

on the pressure of air sealed in a flask at a constant temperature.



• Rubber cork

• 50 ml plastic syringe



• Three-way valve

• Pressure sensor (150 – 1150 mbar)




Two syringe needles (20-gauge) are inserted through the cork, until their tips project

slightly out of the cork (Figure 2).

Figure 2


At the other end of one of the needles, projecting out of the upper side of the cork, a

three-way valve (as used in infusions) is attached via a very short latex tube. A 50 ml

syringe is attached to the latex tube on the other end. A Pressure sensor is

connected to the second needle, through another short latex tube.

Turn the valve until its opening is directed vertically. In this position, air can flow

through the valve from the flask to the surroundings. In order to stop airflow, turn the

valve until its opening reaches a horizontal position. In this position air can flow from

the flask to the syringe.



Data Logger Setup

Sensors:


Rate:

Manual

Samples:

20 samples

Experimental Procedure 1. Fill the glass flask with water. Leave a small volume of air in the flask, to prevent

water penetration into the needles.

2. Close tightly the flask with the cork. Make sure that the system is completely sealed

from air penetration.


4. Follow the Pressure level in the flask. Collect the data manually. Click Run on

the upper toolbar each time you wish to record a data sample.


5. Turn the valve attached to the cork of the flask, until an atmospheric pressure (about

1000 mbar) is obtained in the flask. Turn the valve again to enable airflow from the

flask to the syringe.

6. Draw the syringe’s piston to a volume of about 5 ml. Measure the new value of

pressure obtained after the increase in the volume of air in the system (syringe air

volume added to air volume in the flask).

7. Repeat this step for at least nine more samples. Draw the piston for an additional

volume of 5 ml each time.

8. Follow changes in pressure during the experiment.



Data Analysis 1. Add a column of the volumes' values:

a. Click Table view on the main toolbar to display the table.

b. Click Tools on the menu bar, then click Add Manual column to open the dialog.

c. Type Volume in the Column title text box, type mL in the Column unit text box,

then click OK.

d. MultiLab opens a new column in the table window. Insert the volume data by

tapping on a cell and typing in the appropriate value:

Figure 3


e. Click Graph view to return to the graph view

2. From theory we know that the volume is inversely related to pressure. To verify this

we will produce a graph of the Pressure vs. Volume:

a. Click Format graph on the graph (lower) toolbar.

b. Select Volume in the X-axis list, and then click OK.

c. Remove the Pressure data by selecting it in the Data Map and clicking Hide.

3. Since increase in volume is inversely related to pressure the graph of the Pressure

vs. 1/Volume should be a linear curve. Apply the reciprocal function to the volume

data:

a. Click Analysis Wizard on the main toolbar.

b. Select Reciprocal in the Functions drop-down menu.

c. Select Edited data - Volume in the G1 drop-down menu.

d. Type 1/Volume in the Name text box.

e. Click OK.

4. Produce a graph of the Pressure vs. 1/Volume:

a. Click on the Pressure data on the Data Map and then click Show on the lower

toolbar.

b. Click Format graph on the graph (lower) toolbar.

c. Select 1/Volume in the X-axis drop-down menu, and then click OK.

d. Remove the Pressure data by selecting it on the Data Map and clicking Hide.


The graph of the Pressure vs. Samples obtained in the experiment is shown

below:

Figure 4

Questions 1. By what means was the volume of air in the system changed?

2. Air volume was changed in a stepwise manner and data collected manually. What is

the meaning of each point collected in this way?

3. How is the pressure developed in the experiment related to the changes in air

volume? Support your conclusions with the results you obtained in the experiment.

4. The changes in volume were done by adding a constant value at each step. The

shape of the graph obtained exhibits a 1/x pattern. What is the explanation for that?

Further Suggestions Start the experiment with a 1 ml syringe attached to the cork, and the syringe‘s piston

extended. Compress the air by pushing the piston. Follow the changes in pressure

resulting from the compression.

Hess’ Law: Additivity of Heat of Reactions 59

10. Hess’ Law: Additivity of Heat of Reactions

Figure 1

Introduction According to Hess’s Law, if a reaction can be carried out in a series of steps, the sum of

the enthalpies of the individual steps should equal the enthalpy change for the overall

reaction. The reactions we will be using in this experiment are:

1. Solid sodium hydroxide dissolved in water to form an aqueous solution of ions.

NaOH(s) Na+(aq) + OH-

(aq)

2. Solid sodium hydroxide reacts with aqueous hydrochloric acid to form water

60 Hess’ Law: Additivity of Heat of Reactions

and an aqueous solution of sodium chloride.

NaOH(s) + H+(aq) + Cl-(aq) H2O(l) + Na(aq) + Cl-(aq)

3. Solutions of aqueous sodium hydroxide and hydrochloric acid react to form water

ioned aqueous sodium chloride.

Na+(aq) + OH- (aq) + H+

(aq) + Cl-(aq) H2O(l) + Na+(aq) + Cl-(aq)



• 250 ml beaker



• 50 ml of 1.0 M NaOH

• 50 ml of 1.0 M HCl

• 100 ml of 0.5 M HCl

• 100 ml of water

• 4 g of solid NaOH

• Safety goggles







Data Logger Setup

Sensors:


Rate:

Every second

Samples:

500 samples

Experimental Procedure 1. Obtain and wear goggles.



insert the temperature sensor.

3. Pour 100 ml tap water to the coffee cup.


5. Place the cover on the cup, but leave a narrow opening so that NaOH can be added.




Reaction #1:

9. Add 2 g crystalline NaOH to the cup and immediately replace the cover to ensure

that the cup will be covered.






Reaction #2:

13. Repeat the dissolution reaction (steps 3-12) using 100 ml of 0.5 M HCl instead of

water.

Caution: Handle the HCl and NaOH(s) with great care.

Reaction #3:

14. Repeat steps 3-12, initially measuring out 50 ml of 1.0 M HCl into the beaker. In step

9, instead of solid NaOH, add 50 ml of 1.0 M NaOH.

Data Analysis

1. Use the First cursor and the Second cursor to display the temperature

change, ∆t, for each reaction.









Backward cursor.




Second cursor .



the First cursor.


2. Determine the mass of 100 ml of solution for each reaction (assume the density of

each solution is 1 g/mL).

3. Calculate the heat released by each reaction.

q, (cp=4.18J/g *C) q = Cp x m x ∆t

4. Find ∆H (∆H = -q)

5. Calculate moles of NaOH used in each reaction.

6. Determine ∆H/mol NaOH in each of the three reactions.

7. Combine the heat of reaction (∆H/mol) 1 and 3.

An example of the graph obtained in this experiment is shown below (use the Data Map

to display two reactions on one graph).

Figure 2

Questions 1. According to your data, is the heat of reaction equal to the sum of heats of formation

of the individual components of the reaction?

2. Find the percentage error for the experiment.


Further Suggestions 1. Dissolve different amounts of NaOH. Follow the changes in temperature.

2. Follow the changes in pH in each case.

3. Dissolve a hydrous CuSO4 in water and CuSO4 x 5H2O(s) in water and calculate

Hess’ Law for:

CuSO4(s) + 5H2O CuSO4 x 5H2O

Heat of Combustion 65

11. Heat of Combustion

Figure 1

Introduction In this experiment we will use Hess’s Law of heat reaction that would be difficult to

obtain by direct measurement. We will study the heat of combustion of magnesium

ribbon:

Mg (s) + ½ O2(g) MgO(s)

This equation can be obtained by combining

1. MgO(s) + 2HCl (aq) MgCl2(aq) + H2O

2. Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

66 Heat of Combustion

3. H2(g) + ½ O2(g) H2O(l)



• 250 ml beaker


• Mg(s) – Magnesium ribbon 0.5 g

• 1 g MgO

• 1 M – 500 ml HCl


• Safety goggles






Data Logger Setup

Sensors:


Rate:

Every second

Samples:

500 samples




and larger than the circumference of the coffee cup. Pierce a hole in the cover for the

Temperature sensor.

3. Pour 100 ml 1.0 M HCl to the coffee cup.


5. Put the cover on the cup, but leave a narrow opening so that MgO can be added.


Wait until readings from the sensors are stable.

Reaction #1:

7. Start stirring the HCl in the coffee cup.

Add 1 g crystalline MgO to the cup and immediately replace the cover so that the cup


8. Follow the changes in temperature registered on the monitor, until no further

changes are observed.



Reaction #2:

11. Repeat the reaction (steps 7 - 10) using 0.5 g of magnesium ribbon rather than

magnesium oxide powder.

Data Analysis

1. Use the First cursor and the Second cursor to display the temperature

change, ∆t, for each reaction.










Backward cursor.




Second cursor .



the First cursor.

2. Calculate the heat released by each reaction, q. q =Cp∆t cp = 4.18 J/g*C, m =

100 g of HCl solution.

3. Convert joules to kJ in your final answer.

4. Determine ∆H (∆H = q).

5. Determine the moles of MgO and Mg used.

6. Calculate �H/ mol MgO and Mg.


An example of the graph obtained in this experiment is shown below (use the Data Map

to display the two reactions on one graph):

Figure 2

Questions 1. Determine ∆H/mol Mg for the reaction:

Mg(s) + ½ O2(g) MgO(s)

(∆H for H2(g) + ½ O2(g) H2O(l) = -285.8 kJ)

2. Determine the percentage error for the answer you obtained for Question #1. The

accepted value is 602 KJ.

Further Suggestions Determine the heat of reaction that would result from the following equation:

CuSO4(s) + 5H2O CuSO4 x 5H2O(s)


Saltwater Conductivity 71

12. Saltwater Conductivity

Figure 1

Introduction Dissolving solid sodium chloride in water releases ions according to the following

equation:

NaCl(s) Na+(aq) + Cl-(aq)

In this experiment we will study the effect of increasing sodium chloride concentration on

conductivity. We will be measuring conductivity as the ion concentration of the solution

being monitored is gradually being increased by the addition of concentrated NaCl

drops.

72 Saltwater Conductivity


• 200 ml distilled water

• Conductivity sensor (0-20 mS)


• 50 ml Sodium Chloride solution


• Safety goggles


2. Connect the Conductivity sensor to Input 1 (I/O-1) of the Nova5000.




Data Logger Setup

Sensors:


Rate:

Manual

Samples:

20 samples

Saltwater Conductivity 73


2. Pour 40 ml of distilled water into the flask and place it on a magnetic stirrer.

3. Insert the Conductivity electrode into the water in the flask.

4. Click Run on the upper toolbar to begin recording to measure the conductivity

of the solution in the flask.

5. Collect the data manually: Click Run on the upper toolbar each time you wish

to record a data sample.

6. Start stirring the water in the flask.

7. Add one drop of salt solution to the distilled water. Stir to ensure thorough mixing.

8. Click Run to measure the conductivity of the solution in the flask.

9. Continue this procedure, adding one-drop portions of salt solution, measuring

conductivity, and entering the total number of drops until a total of 10 drops have

been added.

10. Change the rate of stirring during the experiment and follow its effect on the rate of

change in conductance.


12. Save your data by clicking Save on the upper toolbar

Data Analysis Calculate the rate of change in conductance obtained in the process:

1. Use the cursors and to select the desired range.

2. Click Linear fit on the main toolbar. The fit equation will be displayed in the


3. The slope of the fit line is the rate of change in conductance.

74 Saltwater Conductivity

An example of the graph of conductivity and conductance changes, obtained in this

experiment, is shown below. The linear line is the linear fit.

Figure 2

Questions 1. Describe the change in conductivity as the concentration of NaCl solution was

increased by the addition of NaCl drops.

2. What kind of mathematical relationship does there appear to be between conductivity

and concentration?

3. What was the effect of changing the stirring rate on the conductance? Explain this

phenomenon.

Further Suggestions Add a solution of CaCl2 (with the same concentration) in place of NaCl.

The Lambert-Beer Law 75

13. The Lambert-Beer Law

Figure 1

Introduction The amount of light penetrating a solution is known as transmittance, expressed as the

ratio between the intensity of the transmitted light, It, and the initial light intensity of the

light beam, I0:

0II

T t=

76 The Lambert-Beer Law

Where: T – Transmittance

It – Intensity of the transmitted light

I0 – Intensity of the initial light beam

Even though the relationship between transmittance and absorbance would appear to be

a simple inverse relationship, the true relationship between these two variables is

inverse and logarithmic (base 10):

TT

A log)1log( −==

Where: A – Absorption

T – Transmittance as a number between 0 and 1 )100

%(T

The absorption A of a dissolved substance is a linear function of its concentration, the

so-called Lambert-Beer Law. The length of the light path (thickness of the cell) and the

Extinction coefficient (a substance specific constant) determine the slope of the linear

plot.

dcA ε=

Where: c – Concentration

d – Thickness of the cell

ε – Extinction coefficient

The Lambert-Beer Law however is valid only for diluted solutions. The limits for its

validity differ for different materials. As a general rule, one can understand that every

material showing absorption of up to 0.5 - 0.6 still obeys the Lambert-Beer Law.

In this experiment, we follow the behavior of the linear absorption increase of CuSO4

with increasing concentration in order to determine its extinction coefficient at a given

wavelength within its absorption area.



• 4 CuSO4 × 5 H2O - solutions ranging from 0.006 to 0.05 mol/L

• Colorimeter sensor


2. Connect the Colorimeter sensor to Input 1 (I/O-1) of the Nova5000.



Data Logger Setup

Sensors:

Input 1: Colorimeter

Rate:

Manual

Samples:

10 samples

Experimental Procedure 1. Insert the red filter slide into the Colorimeter.

2. Calibrate the Colorimeter sensor (refer to the appropriate chapter in the Colorimeter

sensor sheet).



4. Measure your samples in the order of increasing concentration. Collect the data

manually. Click Run on the upper toolbar each time you wish to record a data

sample.




Figure 2

Data Analysis 1. Note that your initially measured variable is transmission, which must be converted

into absorption:

a. Use the First cursor to select the graph.

b. Click Functions on the main toolbar, and then click the Functions tab.

c. In the Functions drop-down menu select Log10.

d. In the A text box enter -1. In the B text box enter 0.01.

e. In the Name text box enter a name (e.g. Absorption).


f. Click OK.

2. Enter the CuSO4 concentration data into a manual column:


b. Click Tools on the menu bar.

c. Click Add Manual Column.

d. In the Column title text box enter a name (e.g. Concentration).

e. In the Unit text box enter a unit (e.g. ml/L).

f. Click OK.

g. In the table, click the first cell of the new column and enter the concentration of

the first sample. Move to other cells and fill in the column.

3. Display a graph of Absorption vs. Concentration.

a. Hide the data obtained in this experiment (Transmission vs. Sample) by clicking

on the experiment on the Data Map and clicking Hide.

b. Click Format graph on the lower toolbar.

c. Select Concentration in the X-axis drop-down menu and then click OK.

Figure 3


4. Apply linear fit to obtain the extinction coefficient of CuSO4:

a. Click Linear fit on the main toolbar. The fit equation will be displayed in the


b. Divide the slope by the thickness of the cell to obtain the extinction coefficient of

CuSO4.

The Ideal Gas Law 81

14. The Ideal Gas Law

Figure 1

Introduction The Ideal Gas Law is based on the simplification that interactions between the atoms

and the volume of each gas atom or molecule are negligible. This approximation is

permitted for low gas concentrations within a closed system. The so-called Ideal Gas

Equation (developed by Boyle-Marriott and Gay-Lussac) quantifies the relationship

between volume, pressure and temperature.

82 The Ideal Gas Law

pV = nRT p = Pressure (T constant)

T = Temperature (p constant)

V = Volume

n = Number of moles of gas

In this experiment, we investigate the connection between pressure and volume of air

sealed in a flask.



• Rubber stopper

• A 20 ml syringe

• A 20-gauge syringe needle

• Short latex tube

• Long latex tube

• Pressure sensor (0 – 700 kPa)




The syringe is connected to the flask by a three-way valve while the Pressure

sensor is connected directly to the flask.




Data Logger Setup

Sensors:

Input 1: Pressure (0 – 700 kPa)

Rate:

Manual

Samples:

20 samples

Experimental Procedure 1. Close the flask tightly with the rubber stopper and insert the needles connecting the

syringe and the Pressure sensor with the gas space inside the flask.

2. Turn the three-way valve so that the system is closed.


4. Pull the syringe stamp in about six 2 ml increments. Collect the data manually. Click

Run on the upper toolbar each time you wish to record a data sample. Write

the volume increments down, as you will add them later to your data.




Data Analysis The experiment described above results in the following plot of pressure decrease

versus number of collected samples:

Figure 2

1. In order to plot a function presenting the relation between pressure and volume, you

have to add the values for the volume increments to the worksheet. Enter the volume

data to a manual column:


b. Click Tools on the menu bar, then click Add Manual column to open the dialog.

c. Type Volume in the Column title text box. Then type mL in the Column unit text

box, and then click OK.

d. MultiLab opens a new column in the table window. Insert the volume data by

selecting a cell and typing in the appropriate value:


Figure 3

e. Click Graph view to return to the graph view.

2. Since increase in volume is inversely related to pressure the graph of the Pressure

vs. 1/Volume should be a linear curve. Apply the reciprocal function to the volume

data:

a. Click Analysis Wizard on the main toolbar.

b. Select Reciprocal in the Functions drop-down menu.

c. Select Edited data - Volume in the G1 drop-down menu.

d. Type 1/Volume in the Name text box.

e. Click OK.

3. Produce a graph of the Pressure vs. 1/Volume:

a. Click Format graph on the graph lower toolbar.

b. Select 1/Volume in the X-axis drop-down menu, and then click OK.

c. Remove the Pressure data by selecting it on the Data Map and clicking Hide.



Figure 4

4. Apply a linear fit to the graph:

a. Click Linear fit on the main toolbar. The fit equation will be displayed in the


b. The slope of the fit line is the product of the number of moles, the temperature

and the gas constant.

Exploring a Candle Flame 87

15. Exploring a Candle Flame

Figure 1

Introduction A flame does not show everywhere the same temperature at every region. The

characterization of such a flame is possible with a thermocouple (a highly sensitive, fast

responding Temperature sensor ranging from –200 °C to well over 1000 °C). In the

following experiment, we are going to explore a candle flame in three characteristic

areas labeled with 1, 2 and 3 in Figure 1.

88 Exploring a Candle Flame


• Candle

• Matches

• Thermocouple Type K (0 °C to 1250 °C)


2. Connect the Thermocouple sensor to Input 1 (I/O-1) of the Nova5000.



Data Logger Setup

Sensors:

Input 1: Temperature (0 to 1250 °C)

Rate:

10 samples per second

Samples:

1000 samples

Experimental Procedure 1. Light a candle and let it burn for approximately two minutes before starting the

experiment.


3. Record the room temperature before you approach the sensor to the flame (note that

the thermocouple is very sensitive and fast responding).

Exploring a Candle Flame 89

4. Monitor the temperature in three different areas of the flame (see Figure 1):

a. The top of the yellow area 1.

b. The middle of the yellow area 2.

c. The top of the dark area around the wick 3.



Data Analysis Plotting Temperature vs. Time shows us that the flame is hottest at the top and gets

cooler the lower you drag the sensor:

Figure 2

Questions 1. What are the temperatures in each of the flame zones? (1, 2 & 3).

2. Why do you think the hottest part of the flame is at the top?

90 Exploring a Candle Flame

Freezing and Melting of Water 91

16. Freezing and Melting of Water

Figure 1

Introduction Freezing is the process of matter turning from the liquid state into the solid state. Melting

is the process of solids turning into the liquid state. This occurs at the so-called freezing

or melting temperature respectively. In this experiment, we investigate the freezing and

melting temperatures of water according to obtained graphs and compare them to one

another.

92 Freezing and Melting of Water


• Stand with utility clamp

• Graduated cylinder

• Beaker (250 ml)

• Test tube

• Glass rod for stirring

• Water

• Ice cubes

• Salt







Freezing and Melting of Water 93

Data Logger Setup

Sensors:


Rate:

Every 10 second

Samples:

200 samples

Experimental Procedure (Part I, Freezing) 1. Fill ice cubes and cold water into a 250 ml beaker (≥150ml).

2. Fill 1 ml of water into a 5 ml test tube and fix it to the stand with a clamp. You have to

adjust the test tube so that the sample is below the ice water level.

3. Place a Temperature sensor into the water inside the test tube.


5. Add about five spoons of salt to the beaker and stir with a glass rod upon lowering

the test tube into the ice water.

6. Stop recording after about 15 minutes and keep the test tube in the ice-water bath.


Experimental Procedure (Part II, Melting)

1. Click Run on the upper toolbar to begin again recording data, raise the test

tube and fix it into a position above the ice-water bath.

2. Dispose of the ice water and add 150 ml warm tap water to the beaker. Lower the

test tube into the warm water.

3. Stop recording data after another 15 minutes.

94 Freezing and Melting of Water


Data Analysis and Questions Examples of the graphs obtained in the experiments are shown below:

Part I: Freezing water Part II: Melting the frozen water

1. What happened to the water temperature during the freezing stage and during the

melting stage?

2. What is the freezing point and the melting point of the water?

3. How does the freezing point of water compare to the melting point?

4. What is the behavior of the kinetic energy of the water within the test tube? State, if it

increases, decreases or remains the same, when:

a. The temperature changes at the beginning and the end of Part I.

b. The temperature remains constant in Part I.

c. The temperature changes at the beginning and the end of Part II.

d. The temperature remains constant in Part II.

5. Did the potential energy increase or decrease in those parts of Question 4, where

kinetic energy remained constant?

Another Look at Freezing Temperature 95

17. Another Look at Freezing Temperature

Figure 1

Introduction In Experiment 15: Freezing and Melting of Water, we have seen that the temperature of

pure water remains constant at its freezing temperature during the freezing and melting

process. Now we are going to observe the freezing process of phenyl salicylate.

Furthermore, we will observe the effect on the freezing point, when a small amount of

another substance (e.g. benzoic acid) is dissolved in the phenyl salicylate.

96 Another Look at Freezing Temperature


• Three Temperature sensors (-25 °C to 110 °C)

• Beaker (> 1 liter)

• Test tube with 2 g phenyl salicylate

• Test tube with 2 g phenyl salicylate and a few milligrams of benzoic acid

• Glass rod for stirring


2. Connect the Temperature sensors to Input 1 (I/O-1), Input 2 (I/O-2) and Input 3 (I/O-

3) of the Nova5000.




Data Logger Setup

Sensors:




Rate:

Every second

Samples:

1000 samples

Another Look at Freezing Temperature 97

Experimental Procedure 1. Fill about 1 L of water into a beaker. Adjust it to around 30 °C by mixing hot and cold

tap water. Monitor the temperature of this water during the experiment with one of

the Temperature sensors.

2. To melt your samples, prepare a second beaker of hot water (60 to 80 °C).

3. Place the two sample test tubes into the hot water until you see a clear liquid in both

test tubes.

4. Place a Temperature sensor in each sample test tube and move them quickly to the

beaker whose water was adjusted to 30 °C.


6. Stir your samples slightly as long as they are melted.

7. Pay attention to the sample test tubes, as the clear liquid will soon contain increasing

amounts of white precipitates, until the liquid is completely frozen.


Data Analysis 1. What was the freezing point of pure phenyl salicylate according to your data?

2. What was the freezing point of the benzoic acid-phenyl salicylate mixture?

98 Another Look at Freezing Temperature


Figure 2

Questions 1. Based on this experiment and its result, what method could a chemist use to check

whether an unknown liquid is pure or has an impurity dissolved in it?

2. Why does the freezing point change by dissolving another component in phenyl

salicylic acid?

3. Does the nature of this change depend on the freezing point of the dissolved

compound?

Freezing temperature of pure

phenyl salicylate

Freezing temperature of benzoic

acid-phenyl salicylate mixture

Energy Contents of Foods 99

18. Energy Contents of Foods

Figure 1

Introduction All human activity requires burning food to receive the necessary energy. In this

experiment, we are going to determine the energy released (in kJ/g) by three food

samples (popcorn, marshmallows and peanuts) by burning them. The released energy

heats a known amount of water and can be calculated from the first equation. You obtain

the energy content by dividing the obtained heat with the mass of the burnt food (second

equation):

100 Energy Contents of Foods

Q = Cp ∆T Cp

∆T

Q

=

=

=

Heat capacity of the water

Temperature change of the water

Quantity of heat emitted/absorbed

Efood = Q/mfood Efood

mfood

=

=

Energy content of the food

Mass of the burnt food



• Stand with ring stand

• Lid or a small can (≤ 50 ml) as a food holder

• Small can (100 – 200 ml) as a water can

• Three food samples

• Scale

• Two stirring rods


• Cold water

• Matches







Data Logger Setup

Sensors:


Rate:

Every second

Samples:

200 samples

Experimental Procedure 1. Determine the mass of the empty food and water can(s).

2. Determine the mass of the food sample that you are going to measure.

3. Add 50 ml cold water into the water can and determine its exact mass.

4. Place the first food sample into the food can. Note that the samples are light and

much easier upon grinding them (this applies especially to the peanuts).

5. Place the food can with the sample directly under the water can.

6. Insert the Temperature sensor into the water (it must not touch the bottom).

7. Start stirring the water sample in the water can.


9. Wait approximately one minute before you light the food sample by using a match.

10. Continue stirring the water sample until the temperature stops rising.



13. Repeat procedure 1-12 for the other two food samples.


Examples of the graphs obtained in this experiment are shown below:

Figure 2: Burning of 4.4 g peanuts

Figure 3: Burning of 0.5 g popcorn


Figure 4: Burning of 3.6 g marshmallows

Data Analysis

1. Use the First cursor and the Second cursor to display the change of

temperature in water (∆T) for each food sample.









Backward cursor.




Second cursor .




the First cursor.

2. Calculate the heat (Q) absorbed by the water according to the first equation.

3. Determine the mass of the sample’s left over and calculate the difference,

which we call mfood as this is the sample fraction that got burnt.

Note: The specific heat capacity of water at 25 °C is 4.18 J/g°C.

Questions 1. Which food had the highest energy content (in kJ/g)?

2. Food energy is expressed in a unit called Calories (4.18 kJ). How many Calories

does a bag of 50 g peanuts have?

3. Peanuts have a high fat content. Marshmallows and popcorn have high carbohydrate

content. What generalization can you make from your results about the relative

energy content of fats and carbohydrates?

Chemical Equilibrium: Finding a Constant, Kc 105

19. Chemical Equilibrium: Finding a Constant, Kc

Figure 1

Introduction According to common understanding, chemical reactions occur with complete

conversions of the initial compounds (educts) to the products. However, there are a

large number of reactions resulting in a mixture of educts and products. One example is

the reaction of iron (III) ions (Fe3+) with thiocyanate ions (SCN—) forming

thiocyanoiron(III) (FeSCN2+):

106 Chemical Equilibrium: Finding a Constant, Kc

Fe3+(aq) + SCN-(aq) FeSCN2+(aq)

Kc = [FeSCN2+]eq/[Fe3+]eq×[SCN-]eq

As we see in the equations above, Kc is a measure for the quantity of a formed product

relative to the educts. This means that the higher Kc is, the more product quantity is

formed. If Kc is 1, the educt and product concentrations are equal. To determine Kc, you

have to first determine the concentrations of all the components. As the FeSCN2+

absorbs blue light (and appears therefore with a red color in solution), we are going to

use the blue slide of the Colorimeter sensor to determine its concentration in four

samples compared to a standard solution with a known concentration of FeSCN2+. The

Fe3+ concentration of this standard solution must be about 100 times higher than the one

of the test solutions. This high concentration with a low initial concentration of KSCN

forces the equilibrium far to the right, therefore we can understand that all the SCN- ions

are transferred to FeSCN2+, which equals the initial concentration of the added SCN-.

Remember that the absorption of FeSCN2+ depends linearly on the concentration

according to Lambert-Beer’s Law. From the determined concentrations, we are then able

to calculate the concentrations of the other components (see Data Analysis below).


• Colorimeter sensor

• Plastic cuvettes

• Test tubes

• Thermometer

• Scale

• 0.002 M KSCN

• 0.002 M Fe(NO3)3 in 1 M HNO3

• 0.2 M Fe(NO3)3 in 1 M HNO3 (standard solution)


• Pipettes

• Three 100 ml beakers

Attention!

Fe(NO3)3 solutions are prepared from 1 M HNO3. Handle them with great care.

Equipment Setup Procedure 1. Arrange for four test tubes fitting at least 10 ml of liquid. Label them 1-4.

2. Pour about 30 ml of 0.002 M Fe(NO3)3 solution into a 100 ml beaker and transfer five

ml of it into each of the labeled test tubes.

3. Pour about 25 ml of the 0.002 M KSCN solution into another 100 ml beaker and

transfer 2, 3, 4 and 5 ml of this solution into test tubes 1 to 4 according to the table

below.

4. Dilute each solution in the four test tubes with water according to the table below.

Mix each solution thoroughly with a glass rod. Make sure that you wash and dry the

rod each time before dipping it into the next solution.

Test Tube 0.002 M Fe(NO3)3 (ml) 0.002 M KSCN (ml) H2O (ml)

1 5 2 3

2 5 3 2

3 5 4 1

4 5 5 0

5. Launch MultiLab.

6. Connect the Colorimeter sensor to Input 1 (I/O-1) of the Nova5000.



8. Measure and record the temperature of one of the prepared solutions as the

determined Kc depends on the temperature.


9. Prepare the standard solution of FeSCN2+ by transferring 9 ml of 0.2 M Fe(NO3)3

into a fifth test tube (labeled 5). Transfer 1 ml of 0.002 M KSCN into the same test

tube. Stir thoroughly.

Data Logger Setup

Sensors:

Input 1: Colorimeter (0 - 100%)

Rate:

Manual

Samples:

10 samples

Experimental Procedure 1. Insert the blue filter slide into the Colorimeter.

2. Calibrate the Colorimeter (refer to the Colorimeter sensor sheet).


4. Measure your samples and the reference point. Collect the data manually. Click Run

on the upper toolbar each time you wish to record a data sample. Note that you

initially measure transmission (%T).



7. Determine the corresponding absorption (A) from the measured transmission as

explained in Experiment 13: The Lambert Beer Law.


Data Analysis 1. Prepare a table to write down the results for the calculations below for each of the

four experiments.

2. First calculate the initial concentrations of Fe3+ and SCN-, based on composition and

dilution with water. Refer to step 2 of the Equipment Setup Procedure for the volume

of each component used in experiment 1-4. Note that you need to know the initial

concentrations to proceed with calculating the equilibrium concentrations.

[Fe3+]i = (VFe(NO3)3/Vtotal) × 0.002 M i = Initial

[Fe3+]i must be the same for all the four test tubes, as you added always the same

volume of 0.002 M Fe(NO3)3 solution.

[SCN-]i = (VKSCN/Vtotal) × 0.002 M

3. Calculate [FeSCN2+]eq as follows ([FeSCN2+]std is 0.0002M):

[FeSCN2+]eq = (Aeq/Astd) × [FeSCN2+]std A = Absorption

eq = Equilibrium

std = Standard

4. Calculate the equilibrium concentration for Fe3+ and SCN- ([Fe3+]eq and [SCN-]eq).

[Fe3+]eq = [Fe3+]i - [FeSCN2+]eq

[SCN-]eq = [SCN-]i - [FeSCN2+]eq

5. Calculate Kc from the following equation (see also Introduction):

Kc = [FeSCN2+]eq/[Fe3+]eq [SCN-]eq

6. Determine an average value for Kc. How constant were your values?


An example of the Absorption vs. Test Tube graph obtained in this experiment is shown

below:

Figure 2

Questions 1. What does Kc express?

2. What happens to Kc, if you add an additional aliquot of one of the educt components

(Fe3+ or SCN-)?

3. Why do you have to prepare the Fe(NO3)3 solutions in 1M HNO3 instead of water?

4. The standard concentration of [FeSCN2+] is given to you in this experiment. You

could calculate it also from the measured absorption of the standard solution (test

tube 5). Which additional information do you need in order to calculate it?

Energy Content of Fuels 111

20. Energy Content of Fuels

Figure 1

Introduction In this experiment, you will find and compare the heat of combustion of two different

fuels: Paraffin wax and methanol. Paraffin is a member of a group of compounds called

alkanes being composed. Important alkanes being applied as fuels are gasoline and

diesel oil.

Methanol and Ethanol are used as gasoline additives and gasoline substitutes.

In this experiment, we are going to compare the energy content of paraffin and methanol

by measuring their heat of combustion.

To determine the heat of combustion, we first calculate the energy used to heat a known

amount of water by burning paraffin and methanol with the following formula:

112 Energy Content of Fuels

Q = Cp m ∆T Q=

m =

∆T =

Heat

Mass of water

Temperature change of the water



• Stand with ring stand

• Beaker or a small can (250 ml)

• Scale


• Cold water

• Stirring rod

• Candle

• Methanol burner (e.g. from a Fondue set burner)

• Matches







Data Logger Setup

Sensors:


Rate:

Every second

Samples:

200 samples

Experimental Procedure 1. Determine the mass of the empty water vessel.

2. Add 100 ml cold water into the water can and determine its exact mass.

3. Determine the mass of the candle before starting the experiment.

4. Place the filled water vessel on a ring stand.

5. Insert the Temperature sensor into the water (it must not touch the bottom).

6. Start stirring the water sample in the water can.


8. Wait approximately 30 seconds before lighting the candle.

9. Continue stirring the water sample while it heats up.

10. Extinguish the flame when the water reaches a temperature of 40 °C.

11. Stop logging after the temperature stops rising.


13. Determine the mass of the candle after the experiment (include all the wax-

drippings).

14. Repeat procedure 1-12 for the other two food samples.

15. Replace the candle with a methanol burner and repeat the experiment with 200 ml of

water.

114 Energy Content of Fuels

16. Don’t forget to determine the mass of the methanol burner upon filling it.

17. Cover the burner with a piece of metal upon distinguishing it and let it cool down to

room temperature.

18. Determine the mass of the burner with the left over methanol.

Data Analysis Examples of the graphs obtained in this experiment are shown below:

Figure 2

1. Determine the mass of the water heated.

2. Use the First cursor and the Second cursor to display the change of

temperature in water (∆T).









Backward cursor.





Second cursor .



the First cursor.

3. Calculate the heat (Q) absorbed by the water according to the given equation.

4. Determine the mass of burnt paraffin and methanol.

5. Calculate the %-efficiency for both experiments. Divide your experimental value

[kJ/g] by the appropriate literature values and multiply the answer with 100. The

literature values are 41.5 kJ/g (paraffin) and 30.0 kJ/g (ethanol).

Questions 1. Which fuel produces more energy per gram burned? Give an explanation for the

difference (hint: methanol CH3OH is an oxygenated molecule, paraffin C25H52 does

not contain oxygen).

2. Suggest some advantages of using ethanol (or paraffin) as a fuel.

3. Discuss the heat loss factors contributing to the inefficiency of the experiment.

Documents

Chemistry Experiments for Nova5000