Students 2 Science Thermochemistry 1.9.17

© 2017 Students 2 Science Inc. Revised 1.9.17 All Rights Reserved.

STUDENTS 2 SCIENCE Virtual lab Experiment

Thermochemistry

An investigation of thermochemistry: exo-‐ and endothermic reactions.

A classroom Experiment in Kit Form for Grades 9-‐12

Brief Background:

This experiment is intended to teach students about the basics of thermodynamics. They will learn about exo-‐ and endothermic processes and how energy changes can be additive in a controlled environment. They will be introduced to the first law of thermodynamics and the concept of enthalpy. Students will measure temperature changes associated with dissolution and chemical reactions and make use of these observations to make predictions when they change the reaction conditions. Applications of thermochemical principals to modern uses are demonstrated.

Safety

Students and teachers must wear properly fitting goggles as they prepare for, conduct, and clean up from the activities in the kit. Read and follow all safety warnings. Also review the Materials Safety Data Sheets. Students must wash their hands with soap and water after the activities. The activities described in this kit are intended for students under the direct supervision of teachers.

Kit Contents:

Designed for 13 groups of 2 students working in pairs

Items in each 2-‐student Ziploc bag • vial labeled “Citric Acid” approximately 1/3 full • vial labeled “Sodium Bicarbonate” approximately 1/3 full (grind all clumps

from sodium bicarbonate before filling) • vial labeled “CaCl2” approximately 1/3 full of calcium chloride


• vial labeled “NH4Cl” approximately 1/3 full of ammonium chloride • (2) 9oz cups (water and waste) • (6) 2oz flat cups • tall 2oz cup • measuring cup • pipet • thermometer • (4) small scoops • (1) Hot Hand (half the package) • (1) sodium acetate hand warmer (removed from plastic packaging)

Items to be handed out with each Ziploc bag • “Thermo layout.docx” sheet • “Thermo Worksheet.docx” sheet • S2S Thermochemistry Student Instructions

Additional items supplied to be able to run the lab two more times:

• 1 Bottle Citric Acid (250 g) • 1 Box Sodium Bicarbonate • 1 Bottle Calcium Chloride, CaCl2 (100 g) • 1 Bottle Ammonium Chloride , NH4Cl (100 g) • 26 Hot Hands (half packages to be used)

Please rinse, dry and save any reusable supplies for subsequent use.

Teacher Supplied Items:

Water, Paper towels, Pen or Pencil

Pre-‐lab Set up:

Prior to connecting to the S2S website, provide each 2-‐student team with the experiment contained in the Ziploc bag and 3 worksheets. Remove one of the large 9oz cups from each bag and fill 3/4 with water. Take everything else out of the zip lock bag and set it up as indicated on the provided layout sheet. Leave the 6 small, flat cups and the measuring cup off to the side and place the thermometer in the tall, small cup. You should have paper towels available for cleanup of spills. Each student must have a pen or pencil. All students and instructors should have safety glasses and gloves.


Instructor/Teacher Procedure and guide

1. Thermodynamics A. First Law of Thermodynamics-‐Law of Conservation of Energy. B. Total energy within a closed system will remain constant C. Energy cannot be created or destroyed; it can however be converted

into something else.

2. Endothermic Reactions A. An endothermic process or reaction is one where heat is absorbed

from the system and used by the system to do something. B. When heat is used, the temperature decreases.

3. Exothermic Reactions A. An exothermic process or reaction is one where heat is released into

the system because of some process happening in the system. B. When heat is released, the temperature increases.

First&Law&of&Thermodynamics&

Also&called&the&Law$of$Conserva-on$Of$Energy$!  total&energy&of&an&isolated&system&is&constant&!  if&system&is&not&isolated,&energy&can&escape&!  energy&can&be&neither&created&or&destroyed&&!  energy&can&be&changed&from&one&form&to&another&&

•  thermal(•  chemical(

•  electrical(•  mechanical(

&

Some(forms(of(energy(are:&

!process!or!reac)on!absorbs!heat!energy!from!the!system!

•  breaking!a!chemical!bond!

•  mel)ng!(solid!to!liquid)!

•  vaporiza)on!(liquid!to!gas)!

When!heat!is!used!by!the!system,!!the!temperature!decreases!

Examples:!


4. Endothermic and Exothermic Examples using water A. When steam at 100 degrees Celsius condenses to form water at 100

degrees Celsius, energy is released to the system. B. When water at 0 degrees Celsius freezes to form ice at 0 degrees

Celsius, energy is released. C. When water at 100 degrees Celsius vaporizes to form steam at 100

degrees Celsius, heat is absorbed. D. When ice at 0 degrees Celsius melts to form water at 0 degrees

Celsius, energy is absorbed.

5. Experiment 1: Hot Hands A. The exothermic process in Hot Hands is based on a chemical reaction

where finely powdered iron metal is oxidized by atmospheric oxygen and becomes iron oxide, or rust.

B. This is an exothermic reaction that gives off a lot of heat for a long time.

!process!or!reac)on!releases!heat!energy!to!the!system!

•  making!a!chemical!bond!

•  crystalliza)on!(liquid!to!solid)!

•  condensa)on!(gas!to!liquid)!

•  combus)on!reac)ons!(burning!fuel)!

When!heat!is!added!to!the!system,!!the!temperature!increases!

Examples:!

Endothermic,and,Exothermic,Processes,


C. Place the thermometer directly in front of you so you can see the temperature of the system, which in this case is the room where the packaged hand warmer was kept.

D. Open hand warmer and place it on top of the thermometer bulb. E. After a few minutes, note the temperature. Do not let the hand

warmer sit such that the temperature rises above 65 degrees Celsius.

6. Experiment 2: The Reusable Hand Warmer A. Works by the heat that is released during the crystallization of sodium

acetate trihydrate. B. Sodium acetate trihydrate melts at 58 degrees Celsius. If it is heated

above this temperature and cooled slowly it doesn’t easily go back to a solid. It will stay as a liquid.

C. If a single crystal of sodium acetate trihydrate is introduced into the liquid, it spontaneously crystallizes and releases the stored energy as heat.

D. To get this hand warmer to work, squeeze the button in the pouch and watch what happens.

E. Place the hand warmer on the thermometer and see how warm it gets. It will be less than 58 degrees Celsius because the room is removing some of the heat as it is being released.

7. Flameless Heater and Other Applications A. Used to heat over 1 billion meals for the US armed forces. (MRE’s) B. Contains finely powdered magnesium metal, alloyed with a small

amount of iron, and table salt. C. To activate the reaction, a small amount of water is added, and the

boiling point of water is quickly reached as the reaction proceeds. The salt dissolves in the added water and causes each of the iron and magnesium metal pairs to act like tiny batteries. Since the batteries

Exothermic+Processes:++Chemical+Reac3on+and+Crystalliza3on+

•  Oxida3on+of+Iron+Metal+to+Iron+Oxide+(rust)+

2+Fe0+(s)+++O2+(g)+!+2+FeO+(s)+

•  Crystalliza3on+of+Sodium+Acetate+

C2H3NaO2(aq)+!+C2H3NaO2.+3H20+(s)++


are in water, they short circuit, generating enormous heat. This boils the water, which can then heat a food pack placed inside the flameless heater.

8. Dissolution of Salts: Exothermic vs. Endothermic

A. The first thing that happens is that the ionic bonds between the cation, the positively charged ion, and the anion, the negatively charged ion, are broken. This requires energy and is endothermic.

B. Water molecules also need to be separated so they can align with the individual ions. This too requires energy and is endothermic.

C. Then attractive bonds are made between water and each of the ions. This releases energy and is exothermic.

D. Whether the act of dissolving is endothermic or exothermic depends on the amounts of each of these energies.

!are!Flameless!Ra*on!Heaters!! !are!Meals!Ready2to2Eat!

•  invented!in!1990!by!chemical!engineers!and!physicists!!

Flameless!Ra)on!Heaters:!

Meals!Ready3to!Eat:!•  design!engineers!and!nutri*onists!u*lized!

FRH!technology!

•  developed!stable,!nutri*ous,!and!tasty!food!


9. Experiment 3: Calcium Chloride A. Take the calcium chloride vial and place it on the sheet in the circle

that reads “current chemical”. B. Place one of the 6 small flat cups in the space marked “current test

solution”. C. Take the thermometer and place it into the large cup containing

water. D. Take the calcium chloride scoop and place 4 scoops into the empty

“current test solution” cup. E. Now record the temperature of the water using the Fahrenheit scale.

This is the initial temperature. F. Using the measuring cup, measure 7.5 milliliters of water and add it to

the cup with the calcium chloride. Swirl and use the thermometer as a paddle to help dissolve the salt.

G. Read the thermometer and record the new temperature on your worksheet.

H. Subtract the initial from the final temperature and record this difference.

I. Is the reaction exo-‐ or endo-‐thermic? (exothermic-‐temperature should rise)

10. Clean up and prep for next reaction: A. Take the thermometer from the cup and use the pipet to wash fresh

water over it and into the waste container. B. Return the thermometer to the empty cup. C. Empty the reaction cup into the waste and put it in the space labeled

“used reaction cups”. D. Put a new flat cup in the “current test solution” space

11. Experiment 4: Ammonium Chloride

Dissolving)a)Salt)in)Water)can)be)Endothermic)or)Exothermic)

Dissolu8on)involves)breaking)bonds)(endothermic)--

!  between)ca8on)and)anion)(ionic)bonding))!  between)water)molecules)(hydrogen)bonding))

Dissolu8on)involves)making)bonds)(exothermic)-)

!  between)water)and)ca8on)!  between)water)and)anion))


A. Take the ammonium chloride vial and place it on the sheet in the circle that reads “current chemical”.

B. Take the thermometer and place it into the large cup containing water.

C. Take the ammonium chloride scoop and place 4 scoops into the empty “current test solution” cup.

D. Now record the temperature of the water using the Fahrenheit scale. This is the initial temperature.

E. Using the measuring cup, measure 7.5 milliliters of water and add it to the cup with the ammonium chloride. Swirl and use the thermometer as a paddle to help dissolve the salt.

F. Read the thermometer and record the new temperature on your worksheet.

G. Subtract the initial from the final temperature and record this difference.

H. Is the reaction exo-‐ or endo-‐thermic? (endothermic, temperature should fall)




13. Enthalpy A. Enthalpy, which has the symbol H, is the amount of heat used or

released in a system. B. We are concerned with the change in enthalpy, delta H of a process or

reaction. C. If the change in enthalpy is greater than zero, heat is used and the

process is endothermic. D. If the change in enthalpy is less than zero, heat is released and the

process is exothermic. E. Explanation of the dissolution experiments;

i. Calcium Chloride 1. Positive change in enthalpy (endothermic process)

when we broke the hydrogen bonds of the solvent, water.


2. When the water molecules surrounded and bonded with the calcium and chloride ions, there was a large negative change in enthalpy (exothermic process)

3. This negative change for bond making was larger than the total positive change for bond breaking therefore the process is exothermic.

ii. Ammonium Chloride 1. The positive enthalpy change for breaking the

ammonium and chloride ions apart is very high; a lot of energy was used (endothermic process).

2. When the water molecules surrounded and bonded with the ammonium and chloride ions, there was a relatively small negative change in enthalpy (exothermic process).

3. This negative change for bond making was less than the total positive change for bond breaking therefore the process is endothermic

Introducing+ ++(symbol+H)+

Enthalpy++!  is+the+amount+of+heat+used+or+released+within+a+

system+!  usually+expressed+as+the+change+in+enthalpy+from+

the+start+to+the+finish+of+a+process++If####ΔH#>#0,#then#heat#is#used##!##endothermic#If###ΔH#<#0,#then#heat#is#released#!#exothermic#

+

Exothermic+Solu/on+Forma/on+ Endothermic,Solu0on,Forma0on,


14. Experiment 5: Dissolving both Salts-‐Students will design and test their own hypothesis

A. How to make a solution that is neither exo-‐ or endo-‐thermic? B. How many scoops of each salt are required such that no temperature

change occurs? The maximum number of scoops of either salt is 4 C. Students should refer to their results on their worksheet in order to

generate their hypothesis D. Fill in the table on the worksheet with their experimental hypothesis E. Carry out the experiment with the following modifications:

i. Record the initial water temperature ii. Add the desired number of scoops of calcium chloride to the

new flat cup. iii. Add the desired number of scoops of ammonium chloride to

the same flat cup. iv. Add 7.5 milliliters of water to the flat cup and mix with the

thermometer to dissolve the salts. v. Record the final temperature. vi. Subtract the initial temperature from the final temperature and

record your results. 15. Clean up and prep for next reaction:

A. Take the thermometer from the cup and use the pipet to wash fresh water over it and into the waste container.

B. Return the thermometer to the empty cup. C. Empty the reaction cup into the waste and put it in the space labeled


16. Experiment 6: Sodium bicarbonate and citric acid. A. Run the experiment twice as before using sodium bicarbonate the first

time, and citric acid the second time, being certain to rinse the thermometer in between as before.

B. Record results on the worksheet. C. Is the reaction endo-‐ or exo-‐thermic? (Both are endothermic-‐

temperature decreases) D. Clean up and prep reaction set up as before.

17. Experiment 7: Sodium Carbonate and Citric Acid mixed together A. Record the initial water temperature B. Generate and record a hypothesis of temperature change when 4

scoops of each salt are used together. C. Add 4 scoops of citric acid to a new, clean flat cup D. Add 4 scoops of sodium bicarbonate to the same flat cup


E. Add 7.5 milliliters of water to the cup and mix with the thermometer to dissolve the salts.

F. Record the final temperature G. Subtract the initial temperature from the final temperature and

record your results. H. Was the reaction exo-‐ or endo-‐thermic? (endothermic)

i. Dissolution of the salts releases energy ii. Removal of one of the 3 oxygen atoms and hydrogen atom from

the carbon in the bicarbonate ion uses a large amount of energy.

iii. The energy released in forming carbon dioxide and water is much less than this, so the chemical reaction is endothermic.

18. Disposal and Clean-‐up A. All solutions can be poured down the drain. B. Solid chemicals can be re-‐capped to be used again. C. Experimental supplies (cups and pipets and scoops) can be rinsed and

dried if needed for future experiments.

Conclusions and Wrap up:

!Citric!acid!(aq) ! !Sodium!bicarbonate!

+!

+!

+!!Sodium!citrate! !! ! ! ! ! ! ! ! !Water! !Carbon!dioxide!

Reac%on(of(Sodium(Bicarbonate(and(Water(

NaHCO3'(s)'''''→'''''Na+'(aq)''+''HCO3.'(aq)!!

H20'

Reac%on(of(Bicarbonate(and(Acid(

HCO3.''+'H+'''→'''CO2'(g)!+'H20!


Students have explored endo-‐and exo-‐thermic Reactions. They have learned that these reactions are governed by the choices of reactants and the quantity of energy released or absorbed is directly proportional to the amount of reactants used. Proper choices of reaction conditions provides for useful industrial applications-‐hand warmers, MREs, etc.


Student Instructions:

1. Experiment 1: Hot Hands A. Place the thermometer directly in front of you so you can see the

temperature of the system, which in this case is the room where the packaged hand warmer was kept.

B. Open hand warmer and place it on top of the thermometer bulb. C. After a few minutes, note the temperature. Do not let the hand

warmer sit such that the temperature rises above 65 degrees Celsius.

2. Experiment 2: The Reusable Hand Warmer A. To get this hand warmer to work, squeeze the button in the pouch and

watch what happens. B. Place the hand warmer on the thermometer and see how warm it gets.

Note the temperature.

3. Experiment 3: Calcium Chloride A. Take the calcium chloride vial and place it on the sheet in the circle

that reads “current chemical”. B. Place one of the 6 small flat cups in the space marked “current test

solution”. C. Take the thermometer and place it into the large cup containing

water. D. Take the calcium chloride scoop and place 4 scoops into the empty

“current test solution” cup. E. Now record the temperature of the water using the Fahrenheit scale.

This is the initial temperature. F. Using the measuring cup, measure 7.5 milliliters of water and add it to

the cup with the calcium chloride. Swirl and use the thermometer as a paddle to help dissolve the salt.

G. Read the thermometer and record the new temperature on your worksheet.

H. Subtract the initial from the final temperature and record this difference.

I. Is the reaction exo-‐ or endo-‐thermic?


water over it and into the waste container. B. Return the thermometer to the empty cup.


C. Empty the reaction cup into the waste and put it in the space labeled “used reaction cups”.

D. Put a new flat cup in the “current test solution” space

5. Experiment 4: Ammonium Chloride A. Take the ammonium chloride vial and place it on the sheet in the

circle that reads “current chemical”. B. Take the thermometer and place it into the large cup containing

water. C. Take the ammonium chloride scoop and place 4 scoops into the empty

“current test solution” cup. D. Now record the temperature of the water using the Fahrenheit scale.

This is the initial temperature. E. Using the measuring cup, measure 7.5 milliliters of water and add it to

the cup with the ammonium chloride. Swirl and use the thermometer as a paddle to help dissolve the salt.

F. Read the thermometer and record the new temperature on your worksheet.

G. Subtract the initial from the final temperature and record this difference.

H. Is the reaction exo-‐ or endo-‐thermic?




7. Experiment 5: Dissolving both Salts-‐Students will design and test their own

hypothesis A. How to make a solution that is neither exo-‐ or endo-‐thermic? B. How many scoops of each salt are required such that no temperature

change occurs? The maximum number of scoops of either salt is 4 C. Students should refer to their results on their worksheet in order to

generate their hypothesis D. Fill in the table on the worksheet with your experimental hypothesis E. Carry out the experiment with the following modifications:

i. Record the initial water temperature


ii. Add the desired number of scoops of calcium chloride to the new flat cup.

iii. Add the desired number of scoops of ammonium chloride to the same flat cup.

iv. Add 7.5 milliliters of water to the flat cup and mix with the thermometer to dissolve the salts.

v. Record the final temperature. vi. Subtract the initial temperature from the final temperature and

record your results.




9. Experiment 6: Sodium bicarbonate and citric acid.

A. Run the experiment twice as before using sodium bicarbonate the first time, and citric acid the second time, being certain to rinse the thermometer in between as before.

B. Record results on the worksheet. C. Is the reaction endo-‐ or exo-‐thermic? D. Clean up and prep reaction set up as before.

10. Experiment 7: Sodium Bicarbonate and citric acid mixed together

A. Record the initial water temperature B. Generate and record a hypothesis of temperature change when 4

scoops of each salt are used together. C. Add 4 scoops of citric acid to a new, clean flat cup D. Add 4 scoops of sodium bicarbonate to the same flat cup E. Add 7.5 milliliters of water to the cup and mix with the thermometer

to dissolve the salts. F. Record the final temperature G. Subtract the initial temperature from the final temperature and

record your results. H. Was the reaction exo-‐ or endo-‐thermic?


Cleanup:

All solutions can be disposed of by pouring down the laboratory sink with a large volume of water. Unused chemicals can be re-‐capped and saved for future use. Vials, cups and pipets can be rinsed, dried and re-‐used for subsequent experiments. Any disposables should be discarded in the trash.

Slides:

First&Law&of&Thermodynamics&

Also&called&the&Law$of$Conserva-on$Of$Energy$!  total&energy&of&an&isolated&system&is&constant&!  if&system&is&not&isolated,&energy&can&escape&!  energy&can&be&neither&created&or&destroyed&&!  energy&can&be&changed&from&one&form&to&another&&

•  thermal(•  chemical(

•  electrical(•  mechanical(

&

Some(forms(of(energy(are:&

!process!or!reac)on!absorbs!heat!energy!from!the!system!

•  breaking!a!chemical!bond!

•  mel)ng!(solid!to!liquid)!

•  vaporiza)on!(liquid!to!gas)!

When!heat!is!used!by!the!system,!!the!temperature!decreases!

Examples:!


!process!or!reac)on!releases!heat!energy!to!the!system!

•  making!a!chemical!bond!

•  crystalliza)on!(liquid!to!solid)!

•  condensa)on!(gas!to!liquid)!

•  combus)on!reac)ons!(burning!fuel)!

When!heat!is!added!to!the!system,!!the!temperature!increases!

Examples:!

Endothermic,and,Exothermic,Processes,

Exothermic+Processes:++Chemical+Reac3on+and+Crystalliza3on+

•  Oxida3on+of+Iron+Metal+to+Iron+Oxide+(rust)+

2+Fe0+(s)+++O2+(g)+!+2+FeO+(s)+

•  Crystalliza3on+of+Sodium+Acetate+

C2H3NaO2(aq)+!+C2H3NaO2.+3H20+(s)++


!are!Flameless!Ra*on!Heaters!! !are!Meals!Ready2to2Eat!

•  invented!in!1990!by!chemical!engineers!and!physicists!!

Flameless!Ra)on!Heaters:!

Meals!Ready3to!Eat:!•  design!engineers!and!nutri*onists!u*lized!

FRH!technology!

•  developed!stable,!nutri*ous,!and!tasty!food!

Dissolving)a)Salt)in)Water)can)be)Endothermic)or)Exothermic)

Dissolu8on)involves)breaking)bonds)(endothermic)--

!  between)ca8on)and)anion)(ionic)bonding))!  between)water)molecules)(hydrogen)bonding))

Dissolu8on)involves)making)bonds)(exothermic)-)

!  between)water)and)ca8on)!  between)water)and)anion))


Introducing+ ++(symbol+H)+

Enthalpy++!  is+the+amount+of+heat+used+or+released+within+a+

system+!  usually+expressed+as+the+change+in+enthalpy+from+

the+start+to+the+finish+of+a+process++If####ΔH#>#0,#then#heat#is#used##!##endothermic#If###ΔH#<#0,#then#heat#is#released#!#exothermic#

+

Exothermic+Solu/on+Forma/on+

Endothermic,Solu0on,Forma0on,


Acknowledgements: This experiment was developed in collaboration with the American Chemical Society and is derived from experiments contained in “Chemistry-‐Investigating Your World” in the IYC 2011 book. Further information can be found at www.acs.org/iyckit. Their support is gratefully acknowledged.

!Citric!acid!(aq) ! !Sodium!bicarbonate!

+!

+!

+!!Sodium!citrate! !! ! ! ! ! ! ! ! !Water! !Carbon!dioxide!

Reac%on(of(Sodium(Bicarbonate(and(Water(

NaHCO3'(s)'''''→'''''Na+'(aq)''+''HCO3.'(aq)!!

H20'

Reac%on(of(Bicarbonate(and(Acid(

HCO3.''+'H+'''→'''CO2'(g)!+'H20!



Names ___________________________________________

Results Worksheet

Part 1 Results

Test Compound Scoops of Salt

mL of Water

Temperature ( °F )

Initial (TI) Final (TF) Difference (TF – TI)

CaCl2 Calcium Chloride 4 7.5

NH4Cl

Ammonium Chloride 4 7.5

CaCl2 + NH4Cl CaCl2=

NH4Cl= 7.5

Hypothesis


mL of Water

Temperature ( °F )

Initial (TI) Expected Final (TF)

Expected Difference (TF – TI)

CaCl2 Calcium Chloride 7.5

NH4Cl

Ammonium Chloride 7.5

CaCl2 + NH4Cl CaCl2=

NH4Cl= 7.5

Experimental Results : Is Hypothesis Confirmed? (Yes/No) __________ Why?


Part 2 Results


mL of Water

Temperature ( °F )

Initial (TI) Final (TF) Difference (TF – TI)

Sodium Bicarbonate 4 7.5

Citric Acid 4 7.5

Citric Acid + Sodium Bicarbonate

4 each 7.5

Hypothesis


mL of Water

Temperature ( °F )

Initial (TI) Expected Final (TF)

Expected Difference (TF – TI)

Citric Acid + Sodium Bicarbonate 4 each 7.5

Experimental Results : Is Hypothesis Confirmed? (Yes/No) __________ Why?


Documents

Students 2 Science Thermochemistry 1.9.17