cooperative learning in chemistry

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EMS 375 Teaching Portfolio Deborah L. Boxall April 29, 2005

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Page 1: Cooperative Learning in Chemistry

EMS 375 Teaching Portfolio

Deborah L. BoxallApril 29, 2005

Page 2: Cooperative Learning in Chemistry

TABLE OF CONTENTS

Artifacts for student development (descriptions) ...........................................................................3Learning planner .......................................................................................................................4Kinetic theory of gases: student misconceptions ......................................................................5STAD cooperative learning lesson plan ...................................................................................6

Artifacts for multiple instructional strategies (descriptions) ........................................................ 13Direct instruction lesson plan.................................................................................................. 145-E Learning cycle lesson plan ............................................................................................... 21Coin Calorimetry lab worksheet ............................................................................................. 25

Artifacts for motivation and management (descriptions) ............................................................. 29Free Energy worksheet ........................................................................................................... 30Free Energy quiz ..................................................................................................................... 32STAD lesson plan signs .......................................................................................................... 34Team recognition certificates.................................................................................................. 36

Artifacts for communication and technology (descriptions) ........................................................ 38Specific heat slide show (external filename: specificheat.pps)TI-83/CBL-2 Instruction sheet................................................................................................ 39Kinetic Theory of Gases: Technology Integration.................................................................. 40Reaction rates slide show (external filename: reactionrates.pps)Script for reaction rates slide show......................................................................................... 42

Artifacts for planning (descriptions)............................................................................................. 43Kinetic Theory of Gases: Cross-Curricular Integration.......................................................... 44Token economy....................................................................................................................... 46

Artifacts for assessment (descriptions) ......................................................................................... 48Enthalpy of solution worksheet .............................................................................................. 49Kinetic theory of gases: unit test............................................................................................. 53Thermochemistry: unit test (traditional) ................................................................................ 56Thermochemistry: alternative unit test ................................................................................... 64Concept interview ................................................................................................................... 66Graphic organizer.................................................................................................................... 68Transparency 2 (STAD).......................................................................................................... 71

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(Page numbers are linked)
Page 3: Cooperative Learning in Chemistry

ARTIFACTS FOR STUDENT DEVELOPMENT

Instruction to enhance cognitive development: On page 4 is a scanned version of a poster entitled Learning Planner that was prepared toillustrate my philosophy of education for an Educational Psychology class in which I wasenrolled during Spring, 2004. As can be seen from examination of the text on the poster, Ibelieve that students learn best when given the opportunity to examine their own misconceptionsand preconceptions through self-evaluative quizzes. Students must learn a variety of newinformation ranging from factual content knowledge to the more in-depth conceptualunderstanding. While the factual content knowledge is best taught using frequent reviews,engaging the students in active discussion is a more efficacious teaching strategy for the moreconceptual topics. A variety of learning styles are also addressed ranging from auditory (Singalong: The Element Song), kinesthetic (Acting Out) to visual (Re-evaluation: ConceptMapping).

Kinetic Theory of Gases: Student Misconceptions (pg. 5) illustrates some of the preconceptionsthat students may have about this particular topic that could potentially prevent learning. Ateacher can identify such potential misconceptions using a formative informal concept interviewon the topic.

Instruction to enhance social development:On pages 6 through 12 is a Student Teams-Achievement Division (STAD) Cooperative LearningLesson Plan that has been included in this portfolio as an example of a teaching methodologythat promotes improved student social skills as well as introducing students to thermodynamics.The instruction component utilizes directed questioning so that the topic can be developed fromstudent responses. After the central topics (enthalpy, entropy and free energy) have beenintroduced, the students are divided out into groups of three or four where they work together tocomplete the Free Energy Worksheet. While completion of the worksheet is done as a team,individual student accountability is ensured by administering the Free Energy Quiz on theinformation shortly after the worksheet is completed.

Page 4: Cooperative Learning in Chemistry
Page 5: Cooperative Learning in Chemistry

Kinetic Theory of Gases Student Misconceptions

Instead of thinking that: Students may think that: 1. Compression of a gas results in

condensation of the gas into a liquid, which has a definite volume.

1. Gases can be compressed to zero volume.

2. Gas pressure is the amount of force per unit area exerted by the gas molecules on the walls of the container.

2. Gas pressure is the same as force.

3. Increasing the pressure on a gas decreases the average distance between the gas molecules.

3. Increasing the pressure on a gas makes the gas molecules get smaller.

4. Adding heat energy to a gas causes the average intermolecular distance to increase.

4. Heating a gas makes the molecules get bigger.

5. An expanding gas has to work against the pressure exerted by the atmosphere.

5. A gas expanding against the atmosphere is in free expansion.

6. Pure gases and gas mixtures at the same temperature and pressure occupy the same volume regardless of their composition.

6. Gas mixtures occupy more volume than pure gases at the same temperature and pressure.

7. Gases at the same temperature and volume exert the same pressure. While the higher molecular weight gases will have a lower average velocity than the lighter gases, the momentum of the two types of gases is the same.

7. Low molecular weight gases exert greater pressure on the walls of their container than high molecular weight gases in the same container at the same temperature.

8. The space between gas molecules is empty.

8. The space between gas molecules is filled with air.

9. Since gases occupy the entire volume of their container, releasing some of the gas causes the remaining gas to redistribute within the container.

9. A void is formed within the container when some of the gas in the container is released.

10. Heating air makes it rise because the intermolecular distance increases, which causes the density of the air to decrease.

10. Heating air makes it rise because heat rises.

Page 6: Cooperative Learning in Chemistry

THERMOCHEMISTRY:Free Energy, Enthalpy and Entropy

Academic ChemistryGrades 9-12Textbook: Chemistry: Addison-Wesley, 5th ed., (Prentice Hall: 2000), Chpt. 11 (pp 307-13),Chpt. 19 (pp 551-7).

SWBAT:• Distinguish between physical and chemical changes.• Identify the direction of heat flow from calorimetric data.• Define the terms endothermic, exothermic, free energy, and entropy.• Distinguish between spontaneous and nonspontaneous reactions and processes.• Apply their understanding of the relationship between free energy, enthalpy and entropy to

deduce the direction of change for these parameters from experimental data.

NCSCOS:• Competency Goal 2: The learner will build an understanding of regularities in chemistry.

2.05 Identify the indicators of chemical change (precipitate formation, gas evolution,absorption or release of heat).

• Competency Goal 3: The learner will build an understanding of energy changes in chemistry.3.02 Analyze the law of conservation of energy, energy transformation, and variousforms of energy involved in chemical reactions.3.04 Analyze calorimetric measurement in simple systems and the energy involved inchanges in state.3.05 Analyze the relationship between energy transfer and disorder in the universe.

Materials:12 copies of Free Energy worksheet12 copies of Free Energy Worksheet: Answer Key7 Group ID cards printed out on card stock26 copies of Free Energy quiz (for second class period)1 Superteam certificate1 Greatteam certificate

Directions for Substitute Teachers:Before class begins, arrange the student tables into groups of four. Place one Group ID card inthe center of each pair of tables. Put Transparency 1 on the overhead projector and tell thestudents to find their assigned group’s table as the students enter the classroom.

Reaction equations that need to be written on the board are indicated in outlined boxes asillustrated below:

(NH4)3PO4 (s) → 3NH4+(aq) + PO4

3- (aq)

Answers to questions that are to be addressed by the students (eg. Q1, Q2..) are given in italics.

STAD Cooperative Lesson Plan

Page 7: Cooperative Learning in Chemistry

Anticipatory Set: (~ 5 minutes)

A thought experiment: 10.0 g of solid ammonium phosphate ((NH4)3PO4) is added to 100.0 g ofwater that is initially at a temperature of 25 °C. The temperature of the water decreases as thesalt dissolves and reaches a final temperature of 22.7 °C when all of the salt has dissolved.

Q1. Thinking back to our previous discussions about calorimetry, is the salt the system or thesurroundings? (the system) What are the surroundings? (the water)

Q2. Was heat energy exchanged between the system and the surroundings (yes). How can youtell? (the temperature of the water decreased)

Q3. Is this like the phase changes we’ve been discussing? (No, because it involves breakingchemical bonds)

Instruction: (~ 35 minutes)

Just like with physical changes, chemical changes require that energy be transferred between thesystem, which is the reacting species, and the surrounding solvent. So let’s think about how wecan deduce the direction of energy flow:

Q4. What does a negative ∆T tell you about the direction of heat flow when the salt dissolves?(the heat energy needed to break the ionic bonds so the salt can dissolve is absorbed from thesurrounding water…heat flows from the surroundings into the system).

This is how you’d write the reaction equation for the dissolution of ammonium phosphate inwater (Write this on the board, leaving enough space for a second equation to be written next toit)

(NH4)3PO4 (s) → 3NH4+(aq) + PO4

3- (aq)

Let’s calculate the amount of heat that was used to break the ionic bonds in the salt:

gfm[(NH4)3PO4] = 114.9 g/mol 10.0 g of (NH4)3PO4 → 8.7 x 10-2 moles (NH4)3PO4

∆T = -2.3 °C mass(H20) = 100.0 g

heat absorbed from water = mCp∆T = (100.0g)(4.184 J/g°C)(-2.3°C) = -962 J

molar heat of reaction = +962 J / 0.087 moles = 11 kJ/mol

*Stress that the heat of the reaction is of the same magnitude but of opposite sign.

Page 8: Cooperative Learning in Chemistry

Exothermic example

Consider the following reaction: (Write it on the board next to the first reaction)

3NH3 (aq) + H3PO4 (aq) → 3NH4+(aq) + PO4

3- (aq)

Q5. If the same molar amounts of the reactants as in the first example (8.7 x 10–2 mol NH3 andH3PO4) are added to 100.0 g of water, would you expect to get the same ∆T? (Most students willsay that it should be the same since are producing the same products)

Actually, this is an example of an exothermic reaction in that heat is transferred from thereactants to the surroundings. A ∆T of +13.3°C would be observed if we were to carry out thisreaction.

Define the terms endothermic (heat absorbed during a reaction) and exothermic (heat given offduring a reaction). Write below the dissolution equation ∆H = +11 kJ/mol. Write below theneutralization equation ∆H = -64 kJ/mol. Emphasize that exothermic reactions have negativeenthalpy values while endothermic are positive.

Free Energy: Spontaneous vs. nonspontaneous reactions

Consider the following reaction:

2NaHCO3 (s) → Na2CO3 (s) + H2O (g) + CO2 (g) ∆Hrxn = +129 kJ/mol

The reactant in this case is sodium bicarbonate, more commonly known as baking soda and it’sdecomposing to form sodium carbonate, water and carbon dioxide. Sodium carbonate is used toform soaps and detergents. Note that this reaction is endothermic.

Q6. In our first thought experiment, ammonium phosphate would spontaneously dissolve inwater even though it was an endothermic reaction. Has anyone ever noticed baking sodaspontaneously decomposing to sodium carbonate? (No…doesn’t happen at room temperature)

Q7. Can you think of a way to make this nonspontaneous reaction become spontaneous…tomake it happen? (Heat it up)

The reason why it was necessary to add heat to the second endothermic reaction before it wouldoccur has to do with a new term that I’d like to introduce to you…

∆Grxn = Free energy of the reaction

See page 549 of text for definition: free energy is energy that’s available to do work.

Page 9: Cooperative Learning in Chemistry

Perhaps more significant to our study of chemistry is to state that spontaneous processes releasefree energy and nonspontaneous processes absorb free energy. Hence:

If ∆G < 0 → SPONTANEOUS If ∆G > 0 → NONSPONTANEOUS

Q8: So, thinking back to our exothermic and endothermic reactions (point)…are theyspontaneous reactions or not? (Yes…the temperature change tells us that the reactionoccurred…hence was spontaneous)

Let’s consider use the endothermic reaction to try to tease out the relationship between enthalpyand free energy. We know that it’s a spontaneous reaction so ∆G is negative. We know that it’san endothermic reaction so ∆H is positive. The only way a positive enthalpy value can lead to anegative free energy value is if some value larger than the enthalpy value is subtracted from theenthalpy:

∆G = ∆H – something

Next consideration…a nonspontaneous process (a positive ∆G) with a large endothermicenthalpy value such as the decomposition of sodium bicarbonate can be made spontaneous byincreasing the temperature…that tells us that temperature has something to do with the‘something’ in this expression. (erase ‘something’ and change to the following)

∆G = ∆H – T∆S

This new term, ∆S, is the change in the entropy of the system. Entropy is defined as being thedegree of disorder or randomness of the system. For example, when a solid is melted its entropyincreases a little bit, but not nearly as much as the increase in entropy that’s observed when aliquid is vaporized.

∆S ≈ 0 for solid → liquid ∆S > 0 for liquid → gas

Another example of entropy in action…Q9: What combination of heads and tails would you expect to get the most often if you were totoss 10 pennies 100 times? (50:50 mix of heads and tails) What combination would you expectto get the least often (all heads or all tails)

The reason you expect to get the 50:50 mix the most often is because it can be obtained thegreatest number of ways, while there is only one way to get either all heads or all tails. The50:50 mix of coins has the greatest degree of disorder in it and thus the greatest entropy.

Looking back once again at our endothermic reaction, the reason this reaction occursspontaneously is because the entropy change is large enough to counteract the effect of the

Page 10: Cooperative Learning in Chemistry

positive enthalpy value and result in a negative free energy value. The chemical system has gonefrom one equivalent of a highly organized solid lattice to two equivalents of ions that aredispersed throughout the solution.

Guided Practice: (~ 15 minutes)Put up Transparency 2 and fill in the columns under ∆G, ∆H and ∆S. Assess studentunderstanding by their ability to deduce the signs for ∆G, ∆H and ∆S.

Let’s look at a few more examples:What can we say about enthalpy, entropy and free energy in each of the following?

H2O (s) → H2O (l), T = 293 KQ10. What is 293 K in Celsius? (20°C) Ice will spontaneously melt at 20°C, so –∆G, heat isabsorbed from the surroundings when water undergoes a phase change from solid to liquid so+∆H, and there is small increase in the disorder to the system so we’ll write it as 0/+.

H2O (s) → H2O (l), T = 253 KQ11. What is the temperature now? (-20°C) Ice will not spontaneously melt at –20°C so +∆G,enthalpy is still positive, and the ∆S would still be 0/+.

H2O (l) → H2O (g), T = 383 K (temp = 110°C)Q12. Does water spontaneously go to its gas phase at this temperature? (Yes) What is the sign on∆G? (negative) What is the sign on ∆H? (positive) Since a gas is being formed, ∆S is large andpositive.

NaOH (s) → Na+ (aq) + OH- (aq), ∆T > 0The temperature change tells us that the reaction occurs (-∆G) and that it is exothermic (-∆H).Going from a solid to solvated ions so ∆S is positive, but probably not very large.(0/+).

CsCl (s) → Cs+ (aq) + Cl- (aq), ∆T < 0Q13. What is the sign of ∆G? (negative) How about ∆H (positive…endothermic) and ∆S? (smalland positive).

2C2H2 (g) + 5O2 (g) → 4CO2 (g) + 2H2O (g), ∆T > 0Q14. What is the sign of ∆G? (negative) How about ∆H (negative…exothermic)? What about∆S? (negative…entropy is decreasing because the reaction started out with 7 equivalents of gasphase reactants and formed only 6 equivalents of gas phase products.)

Q15. What would ∆S be if the H2O formed was in its liquid phase? (∆S would be negative andlarger in magnitude since have 7 equivalents of gas phase reactants going to only 4 equivalentsof gas phase products.)

Page 11: Cooperative Learning in Chemistry

OK…let’s review before we break up into our teams: (Put up transparency 3…text below)

• Chemical reactions can either give off or absorb heat energy.• If heat is absorbed, the heat of reaction is endothermic. (∆H > 0)• If heat is given off, the heat of reaction is exothermic. (∆H < 0)• Spontaneous reactions always have a negative free energy.• Spontaneous endothermic reactions are driven by a large positive increase in entropy.• Nonspontaneous reactions (∆G > 0) can be made spontaneous by increasing the temperature

of the reaction.• The relationship between free energy, enthalpy, temperature and entropy is given by Gibb’s

Free Energy expression:

∆G = ∆H – T∆S

********************** End of Direct Instruction Component ***********************

Team Study: (30 minutes total)Put up transparency 1 and tell the students to find their team tables if they haven’t already doneso. Once everyone is seated in their teams, tell them that their first task is to choose a teamname. Remind the students that they’ll be working in these teams for several weeks so they wantto pick something they’ll want to use for that long. After about 5 minutes, hand out theworksheets and answer sheets: 2 each to the groups of four and only one worksheet and answersheet to the teams of three. This will ensure that the students work cooperatively on completingthe worksheets rather than independently. Tell the students that they have 20 minutes to work onthe worksheets together this class period and that they’ll have 20 more minutes next class beforethey take the quiz on the material.

Five minutes before the end of the class period, have the students gather up the materials(worksheets, answer sheets, and Group ID tags) and bring them up to the front of the room.

SECOND CLASS PERIODAs with the first class period, arrange the student tables into groups of four. Place one Group IDcard in the center of each pair of tables. Also put the worksheets and answer sheets out on thetables. Put Transparency 1 on the overhead projector and tell the students to find their assignedgroup’s table as the students enter the classroom.

After the students are seated in their teams, put Transparency 2 back up on the projector and tellthe students they have 20 minutes to study the material together before they take a quiz on it.

Quiz: (40 minutes)

Closure: (30 minutes)After collecting the quizzes, go over the ∆G, ∆H, ∆S table from the quiz to reemphasize therelationships between temperature change, heat flow and the free energy of the reactions. Also

Page 12: Cooperative Learning in Chemistry

emphasize the importance of paying attention to the phases that are involved when evaluatingextent and direction of entropy change.

Use the remainder of the class period to introduce calculations involving ∆H, ∆G and ∆S (nextlesson plan).

THIRD CLASS PERIOD

At the beginning of the class, announce each teams improvement points from the quiz the daybefore. Direct the students’ attention to the two signs posted in the classroom that explain howimprovement points are awarded and the criteria for winning the Greatteam and Superteamawards (signs attached). Present the Superteam and Greatteam Certificates to the winning teams.Have the students pin their certificates up on the bulletin board for that week’s winners.

Page 13: Cooperative Learning in Chemistry

ARTIFACTS FOR MULTIPLE INSTRUCTIONAL STRATEGIES

Multiple instructional strategies:In addition to the STAD cooperative learning lesson plan, lesson plans utilizing direct instruction(pp. 14-20) and the 5-E learning cycle (pp. 21-24) methodologies are also included. Thedetermination of which lesson plan would yield the greatest degree of learning would dependupon the intellectual sophistication of the students. For example, students with relatively weakcritical thinking skills would learn best under direct instruction, while the more advancedstudents would achieve greater intrinsic motivation to learn the material from the more hands-on,discovery-based approach utilized in the 5-E learning cycle.

Problem based learning:The Coin Calorimetry lab worksheet has been included (pp. 25-28) as an example of aninstructional focus that required the students to demonstrate both problem solving andperformance skills. The worksheet contained enough prompts to facilitate the students derivingthe equation relating heat and specific heat for a particular substance, as well as performing thenecessary calorimetric calculations needed to elucidate the specific heat of the coins used in theexperiment.

Page 14: Cooperative Learning in Chemistry

Direct Instruction Lesson Plan Deborah L. BoxallAcademic Chemistry

1

CALCULATING HEATING CURVES:Phase Transitions and Specific Heat

SWBAT:1. Distinguish between heat and temperature2. Evaluate the relative heat content of different masses of the same substance3. Calculate a heating curve of water given the values for specific heat and enthalpy of

transition for the three phases.4. Sketch a cooling curve from a given or calculated heating curve.5. Predict the rate of temperature increase for two different phases of a substance given

the specific heat of the two phases (faster, slower).6. Predict how heating rate will affect the shape of the heating curve.7. Compare (shorter, longer) the duration of two phase transitions given appropriate heat

of transition values.

NCSCOS:Competency Goal 3: The learner will build an understanding of energy changes inchemistry.

3.03 Compare and contrast the nature of heat and temperature.3.04 Analyze calorimetric measurement in simple systems and the energyinvolved in changes in state.

Materials:75 g lead fishing weights75 g wax candlesCalculating a Heating Curve worksheet (2 pages)

Safety:Lead is toxic and should not be ingested. Additionally, the fishing weights should not bethrown as injury may result.

Anticipatory Set: (~ 5 minutes)Show the students a drawing of two beakers sitting on a hot plate labeled as bothcontaining 18 g of water, but one is at –10 °C and the other is at 90 °C. Ask the studentsto make a prediction by a show of hands whether the temperature of the two beakers ofwater will increase by 10°C at the same rate, or at different rates. If students predictdifferent rates, ask them to predict which beaker (the ice or the liquid water) will heat upfaster. Tell them that the object of the lesson is for them to learn how to calculate how asubstance will heat up if supplied by a constant source of external heating.

Instruction: (~15 min)1. Remind students of the fundamental differences between the solid, liquid and gaseous

phases of a substance in terms of each phase’s internal energy [NCSCOS 1.04].2. Define heat: the amount of internal energy in a substance.

Page 15: Cooperative Learning in Chemistry

Direct Instruction Lesson Plan Deborah L. BoxallAcademic Chemistry

2

3. Have students give examples of various sources of internal energy using prompts andcues as necessary. [Intermolecular: NCSCOS 1.07; Kinetic molecular theory:NCSCOS 1.06].

4. Define temperature: a measure of the average molecular energy5. Stress that heat is an extensive property while temperature is an intensive property

using this example: a 100 g block of lead vs. a 10 g block of lead, both at 40 °C(remind them that this is a bit above body temperature) contains ten times as muchheat.

6. Assess student understanding of the distinction between heat and temperature byhanding out the fishing weights and candles of equivalent weight. Ask the studentswhether it will take the same amount of heat to increase the temperature of both thePb and the paraffin by 5°C. Once they have answered, put up the transparency withthe specific heat capacities of Pb and paraffin and define specific heat capacity. Stressthe concept that specific heat allows the heating characteristics of different materialsto be compared without having to account for the extensive nature of heat.

7. Point out that the heat capacity of paraffin is almost 20 times larger than that of Pb,and will thus require almost 20x as much heat to raise the temperature by 5°. (Thisexample also evaluates possible material based misconceptions about heat).

8. If students seem confused by concept, provide paraffin/water example for additionalreinforcement.

9. Define phase transitions: solid -> liquid, liquid -> gas10. Ask students if they think the temperature of a heated substance undergoing a phase

change will increase or stay the same. Point out that all of the heat absorbed by asubstance goes to breaking intermolecular bonds and cannot increase the averagemolecular energy of the substance as a whole. Hence, the temperature is invariant.

11. Define ∆Hfus: a measure of the energy contained in the intermolecular bonds in thesolid phase of the substance on a per mole basis (ice example). Point out that ‘fusion’refers to the melting process.

12. Define ∆Hvap: a measure of the energy contained in the intermolecular bonds in theliquid phase of the substance on a per mole basis (water example).

13. Put up the transparency with the completed Worksheet data table. Direct students tocomplete their own data tables while handing out the worksheets. Model doing thecalculations in steps 1-5 by either writing directly onto transparency version of theblank worksheet, or onto the chalkboard. Show the students how to fill in thecalculated values into the worksheet Data Table.

Guided Practice: (~20 minutes)1. Have the students calculate the remainder of the values in the Worksheet Data Table

using steps 1-5 from the Calculating a Heating Curve worksheet.2. Check the students’ plots for accuracy and demonstration of understanding that:

a) specific heat influences temperature change of a substance upon heatingb) temperature is invariant during a phase transitionc) the duration of a phase transition depends upon the enthalpy of that transition

3. Ask the students what they think the answers are to the Questions to Ponder.a) Assess student understanding of the relationship between a heating and cooling

curve by drawing a cooling curve on the board and having the students identify it.

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Direct Instruction Lesson Plan Deborah L. BoxallAcademic Chemistry

3

b) Assess student understanding of the effect of heating rate on the shape of theheating curve by drawing two superimposed heating curves depicting differentheat rates, and have the students identify which curve corresponds to the greaterheating rate.

Closure: (~5 minutes)Show the students what their heating curves should have looked like. Point out that thelower heat capacity of the solid and gaseous phases compared to the liquid phase isreflected in the steeper slope for the two phases (it heats up more quickly). Point out thatthe temperature is invariant during a phase transition, and that the duration of thetransition increases as the enthalpy of the transition increases. Show the students the lasttransparency with H2PEw values on it as well. Ask them to predict the differencesbetween the heating curves for water and for H2PEw. (looking for duration of phasetransition).

Independent Practice: (Homework)1. Calculate the Temperature vs. time values needed to plot a heating curve for flatulase,

H2PEw (a fictitious compound from the planet Flatulon) using the following data:

H2PEw Phase Specific Heat(J g-1 °C-1) Tm = 30 °C

Solid 10.5 ∆Hfus = 20.2 kJ/mol

Liquid 4.0 Tb = 90 °C

Gas 16.5 ∆Hvap = 10.3 kJ/mol

Temperature range: –20 °C to 120 °C Heating rate: 200 W Mass of H2PEw: 130 g moles of H2PEw: 2.1 moles

Assess student understanding using the same criteria as in the Guided Practice.

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Calculating a Heating Curve

Name: _______________________________ Date: ___________________

H2O Phase Specific Heat(J g-1 °C-1) Tm = 0°C

Solid 2.0 ∆Hfus = 6.01 kJ/mol

Liquid 4.18 Τb = 100 °C

Gas 2.0 ∆Hvap = 40.7 kJ/mol

Mass of H2O (g): Moles of H2O:

Initial T (°C): Final T (°C):

Heating rate (W):

Heating curve calculations:1. Calculate temperature change from initial temperature to first phase transition

(∆T1):

2. Calculate amount of heat (q1) required to produce ∆T1 [q = (mass H2O)(Cp)(∆T)]

3. Calculate length of time required to deliver q1 (t = q/W):

4. Calculate amount of heat (qp) required for first phase change[qp = (moles H2O)(∆Hp)]:

5. Calculate length of time required to deliver qp:

6. Repeat steps 1-3 for time needed for second temperature ramp (t3).

7. Repeat steps 4 and 5 for time required for second phase change (t4).

8. Repeat steps 1-3 for time needed for final temperature change (t5).

9. Plot Temperature vs. time using your calculated values.

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Blank Worksheet
Page 18: Cooperative Learning in Chemistry

Calculating a Heating Curve

Data TableTemp.(°C) ∆T (°C) q (J) Time (s) Total time (s)

------------ ------------ 0 0

Plot your calculated heating curve below. Label time spans corresponding to phasetransitions.

Questions to ponder:

1. What would the cooling curve for this sample of water look like?

2. How would the appearance of the plot change if the water were heated at twice theheating rate? One-half of the heating rate?

Page 19: Cooperative Learning in Chemistry

Calculating a Heating Curve

Name: _______________________________ Date: ___________________

H2O Phase Specific Heat(J g-1 °C-1) Tm = 0°C

Solid 2.0 ∆Hfus = 6.01 kJ/mol

Liquid 4.18 Τb = 100 °C

Gas 2.0 ∆Hvap = 40.7 kJ/mol

Mass of H2O (g): 18 g Moles of H2O: 1.0 mole

Initial T (°C): -20 °C Final T (°C): 120 °C

Heating rate (W): 200 W

Heating curve calculations:1. Calculate temperature change from initial temperature to first phase transition

(∆T1): ∆T1 = Tf – Ti = (0 °C – (-20 °C)) = 20 °C

2. Calculate amount of heat (q1) required to produce ∆T1 [q = (mass H2O)(Cp)(∆T)]

q1 = (18 g)(2.0 Jg-1 °C-1)(20 °C) = 720 J

3. Calculate length of time required to deliver q1 (t = q/W):

T1 = 720 J/200 J s-1

4. Calculate amount of heat (qp) required for first phase change[qp = (moles H2O)(∆Hp)]:

qfus = (1.0 mole H2O)(6010 J/mol H2O) = 6010 J

5. Calculate length of time required to deliver qp:

6. Repeat steps 1-3 for time needed for second temperature ramp (t3).

7. Repeat steps 4 and 5 for time required for second phase change (t4).

8. Repeat steps 1-3 for time needed for final temperature change (t5).

9. Plot Temperature vs. time using your calculated values.

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Worksheet Answer Key
Page 20: Cooperative Learning in Chemistry

Calculating a Heating Curve

Data TableTemp.(°C) ∆T (°C) q (J) Time (s) Total time (s)

-20 ------------ ------------ 0 0

0 20 720 3.6 3.6

0 Solid -> liquid 6010 30.1 33.7

100 100 7524 37.6 71.3

100 Liquid -> gas 40700 203.5 274.8

120 20 720 3.6 278.4

Plot your calculated heating curve below. Label time spans corresponding to phasetransitions.

Heating curve for H2O at 200 W

-40

-20

0

20

40

60

80

100

120

140

0 50 100 150 200 250 300time (s)

Tem

pera

ture

(°C

)

Questions to ponder:

1. What would the cooling curve for this sample of water look like?

2. How would the appearance of the plot change if the water were heated at twice theheating rate? One-half of the heating rate?

Liquid -> gas

Solid -> liquid

Page 21: Cooperative Learning in Chemistry

EXOTHERMIC AND ENDOTHERMIC REACTIONS:Heat of solution and Hess’s Law

Academic ChemistryGrades 9-12Textbook: Chemistry: Addison-Wesley, 5th ed., (Prentice Hall: 2000), Chpt. 11 (pp 311-16).

SWBAT:• Distinguish between exothermic and endothermic processes• Use relative temperature changes to compare the endothermicity of two reactions.• Define lattice energy and understand that breaking the ionic bonds in a crystal lattice is

always an endothermic process.• Understand that formation of a hydration sphere around ions is an exothermic process• Relate ion size and charge to the enthalpy of hydration• Relate ion size and charge to the lattice energy of the salt.• Calculate enthalpy of solvation for a salt given the lattice energy and enthalpy of hydration

for the ions comprising the salt.• Predict the relative lattice energies of different salts using their knowledge of how ion size

and charge affects the lattice energy.• Draw a diagram to illustrate the combination of endothermic and exothermic processes

occurring when a salt dissolves.• Use the diagram to calculate the lattice energy of a salt given calorimetric data and enthalpies

of hydration.

NCSCOS:1.07 Assess the structure of ionic compounds relating bonding and molecular geometry tochemical and physical properties.2.01 Analyze periodic nature of trends in chemical properties and examine the use of thePeriodic Table to predict ionic radii.2.05 Identify the indicators of chemical change, such as the absorption or release of heat.3.02 Analyze the law of conservation of energy, energy transformation, and various forms ofenergy involved in chemical reactions.3.03 Compare and contrast the nature of heat and temperature.3.04 Analyze calorimetric measurement in simple systems.

Materials:Ice Melter (mixture of NaCl, KCl, urea, and a surfactant)Table salt (NaCl)Baking soda (NaHCO3)Drano (mixture of NaOH and Al pieces)Wax paperTooth picksPlastic cupsSmall ziplock baggiesMeasuring spoons (~Tbsp size)

Deborah L. Boxall Learning Cycle Lesson Plan

Page 22: Cooperative Learning in Chemistry

SafetyDrano contains sodium hydroxide (NaOH), which is a caustic. Students must wear goggleswhen working with this material and wash their hands thoroughly after they’re finished. Also,the reaction of Drano with water produces a lot of heat and hydrogen gas, so care should betaken to avoid thermal burns and producing a combustible amount of H2 gas. For this reason,small portions of Drano should be added to water in open plastic cups.

Engage (~10 min)Ask the students if any of them had trouble getting home on January 19th (the day a 1” snowfallgridlocked Raleigh). Tell them about how slippery it was because the cars had packed the snowdown into ice and the salt trucks couldn’t get through to salt the roads. Ask the students if theyknow why the roads would have been less slippery if the salt trucks had been able to get through.Try to get several alternatives so that student preconceptions related to the action of salt on ice,and possible prior knowledge of exothermic reactions can be assessed.

Show the students an empty Ice Melter bag, and then read the ingredients and the descriptionfrom the back. Specifically:

Ace Ice MelterRAPIDLY MELTS ICE AND SNOW AND PREVENTS RE-FREEZINGAce Ice Melter features a special custom blend of superior ice meltingingredients. Together they melt even the most stubborn ice an snow and work toprevent re-freezing.• Melts ice down to 0°F (-18 °C).Active IngredientsKClNaClNH2CONH2 (urea)C7O6H14 (methyl-α-D-glucopyranoside; a surfactant)

Put 1 Tbsp of the water each into two baggies. Tell the students that you’re going to add someIce Melter to one of the baggies and ask them to predict what will happen. Make sure to getinput from several students. Have one of the students hold both baggies in either hand, then add1 Tbsp of Ice Melter to one of the baggies and ask the student to mix the contents. Have thestudent describe their observations (should observe that an endothermic reaction occurred ratherthan the exothermic reaction that they most likely predicted).

Assessment of prior knowledge: Students that were previously been exposed to conceptsinvolving solutions (Chpt 17 in text) and the colligative property of freezing point depression(Chpt 18 in text) may be able to relate the activity of the Ice ‘Melter’ to freezing pointdepression. If this idea is brought up during class discussion, ensure that the students understandthat the Ice Melter stops working below –18°C because there is no longer any liquid waterpresent.

Page 23: Cooperative Learning in Chemistry

Explore (30 minutes)Ask the students to describe what they think happened, and list their explanations on the board.Be sure to provide sufficient cueing that the students describe the properties of ionic solids. Tellthem you want them to find out if dissolving a substance always results in an endothermic heatof solution. Have the students do the two-baggie experiment with sodium chloride and sodiumbicarbonate (both are endothermic).

Next, read the back of the Drano bottle to the students and ask them what the precautions tellthem about its likely reaction with water. Since they should predict an exothermic reaction, tellthem that they need to use their plastic cups for the next part of the experiment. Have thestudents put on their goggles, and then to pour a small portion of Drano out onto their piece ofwax paper. Ask them to observe what the Drano looks like (blue and white solid with bits ofmetal present). Tell them to put some water in one of their cups, and then add the Drano to thewater. They should get gas evolution and notice that the solution gets very hot. Ask them whatthey think the difference between the reactions with NaCl, NaHCO3, and Drano is. Usequestions to guide the students into suggesting that it’s the reaction between the NaOH and themetal bits that’s producing the heat. This will also assess whether the students believe that allspontaneous chemical reactions are exothermic. Suggest to the students that they use thetoothpicks to remove the metal bits from the Drano mixture, and then try the dissolution reactionagain. They should observe that the solution still gets very hot, but that there’s no gas evolution.Ask them to explain the heat source in this case.

Explain (15 minutes)Ask the following questions:1. How should I write the reaction equation out for when we dissolved table salt in water? (Be

sure to get the students to provide the states of matter)2. Is heat a product of this reaction? (No) So, is this an endothermic or exothermic reaction?

(endothermic) So is the sign of the enthalpy of solution positive or negative? (positive) andwe know that energy is absorbed when the salt is dissolved.

3. What processes do you think are going on that would absorb energy? (most will only be ableto come up with ionic bond breaking) Use this definition to define lattice energy. Ask thestudents if they think all lattice energy values will be positive. (Yes…the enthalpy ofcrystallization must be negative for a stable crystal lattice to exist). Write all lattice energiesare > 0 on the board.

4. How would I write the reaction equations for the other two reactions that we looked at? (getphases as before, and ask about sign of ∆Hsoln)

5. Which one of these three reactions absorbed the most energy from the surroundings?(NaHCO3) How can you tell? (seemed to get the coldest)

6. What is the most significant difference between these reactions? Why do you think thedissolution of NaOH is exothermic while the other two were endothermic? Put upTransparency 1 with the Enthalpies of hydration covered.

7. Looking at just the lattice energy values, would you have expected the enthalpy of solution ofNaOH to be endothermic or exothermic? Why?

8. What else do you think could be going on to result in an exothermic enthalpy of solution?(list explanations on the board)

Page 24: Cooperative Learning in Chemistry

9. What does it mean when we write that an ion is in its aqueous phase? (that there are watermolecules associated with the ions)

10. Do you think it requires energy or releases energy when water molecules form a hydrationsphere around an ion? (releases…greater stability implies lower energy)

Conclusion to draw from student responses: Energy is absorbed during the dissolution process.However, not all ∆Hsoln are endothermic. Energy released during hydration of ions in the solutioncan be sufficient to result in an exothermic ∆Hsoln.

Elaborate (25 minutes)Introduce Hess’s Law: The enthalpy change of a reaction is the sum of the enthalpy changes foreach step of the reaction. Uncover the equation on Transparency 1. Ask the students what theadvantage would be to using Hess’s Law (don’t have to do the experiment in order to find outwhether a reaction is going to be exo- or endothermic).

Handout the Enthalpy of Solution Worksheet, and use directed questioning to guide the studentsindividually through the worksheet. Encourage them to look up the table depicting the relativeionic radii (pg 399) when comparing lattice energies and heats of hydration listed in Tables 1 and2.

Evaluate (done while the students are working on the worksheets)Completion of the worksheet, especially the Born-Haber cycle diagram, comprises an authenticassessment of student understanding. In order for this to be an effective vehicle, the studentsmust be provided with just enough guidance to complete the worksheet independently. Potentialconceptual difficulties are present in the calculations concerning the Group II halides.

Closure (10 minutes)Draw the two Born-Haber diagrams on the board using student input. Emphasize that ‘heatreleased’ indicates an exothermic reaction, while ‘heat absorbed’ indicates an endothermicreaction. Ask the students how the diagrams would have changed if the Mg salts had been usedinstead (both have exothermic heats of solution of about 100 kJ/mol).

In preparation for the next summative evaluation, tell the students that they will need to be ableto:• Predict how lattice energy and enthalpy of hydration change with ionic size and charge• Apply Hess’s Law to the calculation of ∆Hsoln, lattice energy or ∆Hhyd if given two out of the

three values• Interpret a Born-Haber diagram

Page 25: Cooperative Learning in Chemistry

COIN CALORIMETRY

Procedure:1. Add 400 mL of hot water to a 600 mL beaker, and heat to boiling on a hot plate.2. Count and weigh the coins, weigh the coffee cup calorimeter (two nested styrofoam cups)

empty, and after adding 75 mL of water to the inner cup. Record your data below.3. Put the calorimeter containing 75 mL of water into a 400 mL beaker for additional stability.4. Put the coins in the plastic sandwich bag, and use the tongs to suspend the coins in the

boiling water (Twater > 70 °C is OK).5. Collect temperature data from the calorimeter water for one minute, then move the

temperature probe to the hot water bath. Collect temperature data from the hot water bath forone more minute.

6. Transfer both the temperature probe and the bag of coins to the calorimeter. 7. Collect temperature data from the calorimeter for two minutes, swirling the calorimeter

gently every 30 seconds. Press the [STO] key on the TI-83 to stop data collection.8. Plot out the data and determine the plateau temperature of the calorimeter water before the

coins were added (Ti) and the plateau temperature of the hot water bath before the coinswere transferred to the calorimeter (Ti, coins).

9. Determine the plateau temperature of the calorimeter water after the coins were added (Tf).

Data:Type and number of coins:

Mass of coins: Mass of empty calorimeter:

Mass of calorimeter and water: Mass of water:

Ti (calorimeter): Tf (calorimeter):

∆T (calorimeter): (Tf – Ti)

Ti, coins: ∆T of coins: (Tf – Ti, coins)

Page 26: Cooperative Learning in Chemistry

Data Analysis:1. Complete the picture of the calorimeter containing water and coins below. Indicate the

direction of heat flow on the diagram.

2. Given that the specific heat of water is 4.18 J/g°C, how much heat was absorbed by the waterin the calorimeter?

3. Write a mathematical expression relating heat (q) to specific heat.

4. How much heat was given up by the coins?

5. What is the specific heat of the coinage metal? (Hint: Combine the answers from questions 2and 3, and solve for the specific heat of the metal).

Page 27: Cooperative Learning in Chemistry

COIN CALORIMETRY

Procedure:10. Add 400 mL of hot water to a 600 mL beaker, and heat to boiling on a hot plate.11. Count and weigh the coins, weigh the coffee cup calorimeter (two nested styrofoam cups)

empty, and after adding 75 mL of water to the inner cup. Record your data below.12. Put the calorimeter containing 75 mL of water into a 400 mL beaker for additional stability.13. Put the coins in the plastic sandwich bag, and use the tongs to suspend the coins in the

boiling water (Twater > 70 °C is OK).14. Collect temperature data from the calorimeter water for one minute, then move the

temperature probe to the hot water bath. Collect temperature data from the hot water bath forone more minute.

15. Transfer both the temperature probe and the bag of coins to the calorimeter. 16. Collect temperature data from the calorimeter for two minutes, swirling the calorimeter

gently every 30 seconds. Press the [STO] key on the TI-83 to stop data collection.17. Plot out the data and determine the plateau temperature of the calorimeter water before the

coins were added (Ti) and the plateau temperature of the hot water bath before the coinswere transferred to the calorimeter (Ti, coins).

18. Determine the plateau temperature of the calorimeter water after the coins were added (Tf).

Data:Type and number of coins: See tables for expected data

Mass of coins: Mass of empty calorimeter:

Mass of calorimeter and water: Mass of water: 50.0 g

Ti (calorimeter): 20 °C Tf (calorimeter):

∆T (calorimeter): (Tf – Ti)

Ti, coins: 100 °C ∆T of coins: (Tf – Ti, coins)

Answer key

Page 28: Cooperative Learning in Chemistry

Data Analysis:1. Complete the picture of the calorimeter containing water and coins below. Indicate the

direction of heat flow on the diagram.

2. Given that the specific heat of water is 4.18 J/g°C, how much heat was absorbed by the waterin the calorimeter?Use dimensional analysis to determine

# and typeof coin

∆T (°C) q (J) # coin ∆T q (J)

7 US pennies 3.1 648 14 US pennies 5.9 12342 C Nickels 1.5 314 5 C Nickels 3.7 7734 C Dimes 1.4 293 8 C Dimes 2.7 564

4 C pennies 1.9 397 8 C pennies 3.6 752

3. Write a mathematical expression relating heat (q) to specific heat.

q = (mass H2O)(specific heat)(∆T)

4. How much heat was given up by the coins?qcoins = -qwater

# and typeof coin

qcoins (J) # and type ofcoin

qcoins (J)

7 US pennies -648 14 US pennies -12342 C Nickels -314 5 C Nickels -7734 C Dimes -293 8 C Dimes -564

4 C pennies -397 8 C pennies -752

5. What is the specific heat of the coinage metal? (Hint: Combine the answers from questions 2and 3, and solve for the specific heat of the metal).

Cu: 0.384 J/g°C Ni: 0.444 J/g°C

qH2O

Page 29: Cooperative Learning in Chemistry

ARTIFACTS FOR MOTIVATION AND MANAGEMENT

Individual motivation and engagement:The Free Energy Quiz (pg. 32), which is administered following the STAD team study sessionencourages individual motivation and engagement because their individual grade will depend onhow well they were able to learn the material as part of the group. By making the group membersinterdependent (the team score depends upon the average number of improvement points in thegroup), each individual member is also motivated to interact with their group members andcomplete the Free Energy Worksheet during the group study session so that they can all learn thematerial and do well.

Group motivation and engagement: Since there are not enough Free Energy Worksheets (pp. 30-31) provided for each student tohave their own copy, the members of each group must work together to complete the worksheet.

The STAD lesson plan signs (pp. 34-35) are to be posted in the classroom in order to engendergreater team effort and cooperation. The team recognition certificates (pp. 36-37) provide theextrinsic motivation that might be needed to get all of the groups working toward a commongoal.

Page 30: Cooperative Learning in Chemistry

Free Energy Worksheet

1. Label each of the following as either a physical change (P) or a chemical change (C) :

_____ CO2 (s) → CO2 (g) _____ C6H12O6 (s) + O2 (g) → CO2 (g) + H2O (g)

_____ 2H2 (g) + O2 (g) → 2H2O (g) _____ H2O (g) → H2O (l)

_____ NaCl (s) → NaCl (l) _____NaCl (s) → Na+ (aq) + Cl- (aq)

_____Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

2. What is an endothermic process?

3. What is an exothermic process?

4. How can you tell whether a reaction is endothermic or exothermic?

5. What is free energy?

6. How is the spontaneity of a reaction related to its free energy?

7. What is entropy?

8. Indicate whether each of the following physical and chemical changes results in a positive(+), negative (-) or negligible (0/+ or 0/-) ∆S value:

_____ CO2 (s) → CO2 (g) _____ C6H12O6 (s) + O2 (g) → CO2 (g) + H2O (g)

_____ 2H2 (g) + O2 (g) → 2H2O (g) _____ H2O (g) → H2O (l)

_____ NaCl (s) → NaCl (l) _____NaCl (s) → Na+ (aq) + Cl- (aq)

_____Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

9. What is the mathematical expression that relates free energy, enthalpy and entropy?

10. Is it true that all spontaneous reactions are exothermic? Explain your answer.

11. For each of the following, indicate the sign of ∆G and ∆H. Also, if it is possible to evaluate∆S, indicate whether ∆S is positive, negative or negligible.

a. Dissolving 10.0 g of a solid in 100 g of water results in a ∆T = +3°

b. The melting of ice at room temperature

c. Condensation of water vapor at 120°C

d. Condensation of water vapor at 80°C

e. Sublimation of dry ice at room temperature

Page 31: Cooperative Learning in Chemistry

Free Energy Worksheet: Answer Key

1. Label each of the following as either a physical change (P) or a chemical change (C) :

P CO2 (s) → CO2 (g) C C6H12O6 (s) + O2 (g) → CO2 (g) + H2O (l)

C 2H2 (g) + O2 (g) → 2H2O (g) P H2O (g) → H2O (l)

P NaCl (s) → NaCl (l) C NaCl (s) → Na+ (aq) + Cl- (aq)

C Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

2. What is an endothermic process? One that absorbs heat from the surroundings3. What is an exothermic process? One that gives off heat to the surroundings4. How can you tell whether a reaction is endothermic or exothermic? Endothermic

reactions have negative ∆T’s, exothermic rxns have positive ∆T’s5. What is free energy? Free energy is energy that’s available to do work6. How is the spontaneity of a reaction related to its free energy? Spontaneous reactions

have negative free energy values7. What is entropy? Entropy is the amount of disorder in a system8. Indicate whether each of the following physical and chemical changes results in a positive

(+), negative (-) or negligible (0) ∆S value:

+ CO2 (s) → CO2 (g) 0 C6H12O6 (s) + 6O2 (g) → 6CO2(g) + 6H2O(l)

− 2H2 (g) + O2 (g) → 2H2O (g) − H2O (g) → H2O (l)

0/+ NaCl (s) → NaCl (l) 0/+ NaCl (s) → Na+ (aq) + Cl- (aq)

0/− Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)

9. What is the mathematical expression that relates free energy, enthalpy and entropy?∆G = ∆H-T∆S10. Is it true that all spontaneous reactions are exothermic? Explain your answer. No,

endothermic reactions can be spontaneous if there is a large +∆S to offset theeffect of the +∆H

11. For each of the following, indicate the sign of ∆G and ∆H. Also, if it is possible to evaluate∆S, indicate whether ∆S is positive, negative or negligible.a. Dissolving 10.0 g of a solid in 100 g of water results in a ∆T = +3° -∆G, -∆H, 0/+ ∆Sb. The melting of ice at room temperature -∆G, +∆H, 0/+ ∆Sc. Condensation of water vapor at 120°C +∆G, -∆H, -∆Sd. Condensation of water vapor at 80°C -∆G, -∆H, -∆Se. Sublimation of dry ice at room temperature -∆G, +∆H, +∆S

Page 32: Cooperative Learning in Chemistry

Free Energy Quiz

_____ 1. What is always true for a spontaneous process?

a) –∆H b) +∆H c) –∆G d) +∆G e) +∆S

_____ 2. What is true for an endothermic reaction?

a) –∆H b) +∆H c) –∆G d) +∆G e) +∆S

_____ 3. What is true for a spontaneous endothermic reaction?

a) –∆H,+∆S b) -∆H, -∆G c) +∆H, +∆G d) +∆H, +∆S e) +∆G, +∆S

_____ 4. What is true for a spontaneous exothermic reaction?

a) –∆H,+∆S b) -∆H, -∆G c) +∆H, +∆G d) +∆H, +∆S e) +∆G, +∆S

5. Identify whether each of the following is a physical (P) or chemical (C) process.

2H2 (g) + O2 (g) → 2H2O (l) Ba2+ (aq) + SO42- (aq) → BaSO4 (s)

Hg (l) → Hg (g) C2H5OH (g) → C2H5OH (l)

NaCl (s) → Na+ (aq) + Cl- (aq) CaCO3 (s) → CaO (s) + CO2 (g)

6. Identify whether entropy is increasing (I) or decreasing (D) in each of the following:

2H2 (g) + O2 (g) → 2H2O (l) Ba2+ (aq) + SO42- (aq) → BaSO4 (s)

Hg (l) → Hg (g) C2H5OH (g) → C2H5OH (l)

NaCl (s) → Na+ (aq) + Cl- (aq) CaCO3 (s) → CaO (s) + CO2 (g)

7. Complete the table below by filling in the appropriate signs for ∆G, ∆H and ∆S.

∆G ∆H ∆S

Freezing of water at 25°C

H2O (s) → H2O (l); T = 25°C

Nonspontaneous, exothermic process

NaCl (s) → NaCl (l); T = 25°C [Tmp(NaCl) = 801°C]

Condensation of ethanol vapor at 25°C; Tbp = 78.3°C

Name: Date:

Page 33: Cooperative Learning in Chemistry

Free Energy Quiz

_____ 1. What is always true for a spontaneous process?

a) –∆H b) +∆H c) –∆G d) +∆G e) +∆S

_____ 2. What is true for an endothermic reaction?

a) –∆H b) +∆H c) –∆G d) +∆G e) +∆S

_____ 3. What is true for a spontaneous endothermic reaction?

a) –∆H,+∆S b) -∆H, -∆G c) +∆H, +∆G d) +∆H, +∆S e) +∆G, +∆S

_____ 4. What is always true for a spontaneous exothermic reaction?

a) –∆H,+∆S b) -∆H, -∆G c) +∆H, +∆G d) +∆H, +∆S e) +∆G, +∆S

5. Identify whether each of the following is a physical (P) or chemical (C) process.

C 2H2 (g) + O2 (g) → 2H2O (l) C Ba2+ (aq) + SO42- (aq) → BaSO4 (s)

P Hg (l) → Hg (g) P C2H5OH (g) → C2H5OH (l)

C NaCl (s) → Na+ (aq) + Cl- (aq) C CaCO3 (s) → CaO (s) + CO2 (g)

6. Identify whether entropy is increasing (I) or decreasing (D) in each of the following:

D 2H2 (g) + O2 (g) → 2H2O (l) D Ba2+ (aq) + SO42- (aq) → BaSO4 (s)

I Hg (l) → Hg (g) D C2H5OH (g) → C2H5OH (l)

I NaCl (s) → Na+ (aq) + Cl- (aq) I CaCO3 (s) → CaO (s) + CO2 (g)

7. Complete the table below by filling in the appropriate signs for ∆G, ∆H and ∆S.

∆G ∆H ∆S

Freezing of water at 25°C + − −

H2O (s) → H2O (l); T = 25°C − + +

A nonspontaneous, exothermic process + − −

NaCl (s) → NaCl (l); T = 25°C [Tmp(NaCl) = 801°C] + + +

Condensation of ethanol vapor at 25°C; Tbp = 78.3°C − − −

Name: Answer Key Date:

C

B

D

B

Page 34: Cooperative Learning in Chemistry

If your quiz score is…. ImprovementPoints

> 10 points below base score 0

10 to 1 point below base score 10

< 10 points above base score 20

> 10 points above base score 30

Perfect paper 30

default
STAD sign 1
Page 35: Cooperative Learning in Chemistry

Team AverageImprovement Points Award

> 15

> 20

> 25

default
STAD sign 2
Page 36: Cooperative Learning in Chemistry

The

are hereby recognized for their

outstanding team achievement in

Chemistry.Date

Page 37: Cooperative Learning in Chemistry

The

are hereby congratulated on being a

in Chemistry.Date

Page 38: Cooperative Learning in Chemistry

ARTIFACTS FOR COMMUNICATION AND TECHNOLOGY

Communication strategies: Two of the weaknesses of direct instruction is that information is most often communicatedverbally and the student plays a relatively passive role in the learning process. I have tried toincorporate questioning and illustrations (pictures on the board) into all of my lesson plans sothat the students will be more actively engaged; however, neither of these approaches fullyaddresses the two weaknesses. For this reason, I have prepared a PowerPoint slideshow(external to this portfolio file, specificheat.pps) that uses visuals to communicate the anticipatoryset and definition of specific heat as well as requiring direct student interaction. This PowerPointshow could be used in tandem with a lesson, to augment a lesson, or placed on-line for studentsthat were unable to attend the lesson.

Enrichment through technology: There are three examples of the use of technology to enrich instruction. The first, the CoinCalorimetry Lab, utilized student TI-83 Plus calculators interfaced with the CBL-2 portable datacollection device outfitted with a temperature probe. Students used the data logging functionprovided by the TI-83/CBL-2 system to record their temperature data before and after adding thehot coins to the water in the calorimetry cup. While alcohol thermometers could also have beenused for this experiment, they did not offer the level of precision to distinguish between thespecific heats of copper and nickel, 0.383 J/g°C and 0.444 J/g°C respectively. The step-by stepinstructions for data logging have been included here on pg. 39.The second example is a calculator-based lab (pp. 40-41) that guides the students through thesteps necessary to calculate and graph the atmospheric thermoclines and determine thedependence of pressure upon density of particles.The third example is another external PowerPoint slideshow (reactionrates.pps) that is intendedto be included as a visual aid to a discussion about the effect of surface area upon reaction rates.The script that accompanies this slide show is included on pg 42.

Page 39: Cooperative Learning in Chemistry

COLLECTING TEMPERATURE DATA WITH THE TI-83/CBL-2 PORTABLELABORATORY

The following steps will take you through the set-up and the use of the CBL 2 with the TI-83Plus calculator to conduct a simple temperature data collection.

1. Insert the calculator into the cradle on top of the CBL 2 unit and click into place.

2. Place one end of the six inch unit to unit link cable into the port on the bottom of thecalculator and the other end into the CBL 2 unit.

3. Plug the temperature probe into the port labeled: CH 1 on the left hand side of the CBL 2unit.

4. Turn the calculator [ON].

5. Reset the memory on your calculator by pressing [2ND] then [MEM] which is the [+] key.Then choose "RESET" which is 7 by scrolling down or clicking the [7] key.

6. Choose 1,"All RAM" then 2, "RESET" on the next screen. A warning will be displayed thatsays, "Resetting RAM erases all data and programs from RAM."

7. A screen will be displayed that says "RAM cleared." Press [ENTER] and the word "DONE"will appear on the screen.

8. Put the calculator in receive mode by pressing [2ND] then [LINK]. Arrow over so that"RECEIVE" is highlighted then press [ENTER].

9. The calculator will display a "Waiting" screen. Press the <TRANSFER> button on the CBL2 unit. The CBL 2 unit detects which calculator is connected and sends the appropriateversion of the DataMate software to the calculator.

10. The calculator will display a "Receiving" screen. When the calculator has received theprogram, "Done" will be displayed on the screen.

11. To run the DataMate program, click the [APPS] button on the calculator and then arrow upor down to highlight the DataMate program. Press [ENTER].

12. A screen is displayed which says, "CHECKING SENSORS." If the calculator has identifiedthe temperature sensor correctly, a screen will be displayed that says, "CH 1: TEMP (C) andthe current reading as determined by the temperature sensor.

13. Start collecting data by pressing [2] which is "START." A real-time graph of the temperaturewill be displayed.

14. After you are finished collecting data, press the [STO] key. Your graph will be displayed onthe screen.

15. Use the arrow keys to move through each data point on the graph. The graph can be saved forretrieval at a later time.

Page 40: Cooperative Learning in Chemistry

Kinetic Theory of Gases Technology Integration Lesson

NC Competency Goals: Chemistry: 1.06 Analyze the basic assumptions of kinetic molecular theory and its applications. 2.02 Analyze the mole concept and Avogadro's number and use them to calculate the number of

gas molecules from mass of the gas. Computer/Technology Skills: 3.1 Select and use appropriate technology tools to efficiently collect, analyze, and display data. 3.3 Use a calculator, scientific calculator, or graphing calculator for problem-solving. Lesson objectives:

• Use dimensional analysis to convert from non-SI pressure units to SI units. • Use Gay-Lussac’s discovery that air composition is invariant with altitude to calculate the

effective molar mass of air. • Use Avogadro’s number to calculate the number of molecules of a gas from moles. • Utilize TI-83 or TI-82 graphing calculators to generate plots of altitude vs. temperature,

pressure vs. temperature and pressure vs. number of molecules from atmospheric data. • Observe that pressure is directly dependent upon the number of molecules present.

Lesson plan: Draw the table below on the board, omitting the numbers in the shaded areas. Table of Atmospheric Data

Altitude (km)

Temperature (°C)

Temperature (K)

Pressure (mm Hg)

Pressure (kPa)

Mass in 1-L sample

(g)

Total Molecules

0 20 293 760 101.3 1.20 2.50E+22 5 -12 261 407 54.2 0.73 1.52E+22 10 -45 228 218 29.1 0.41 8.54E+21 12 -60 213 170 22.7 0.37 7.71E+21 20 -53 220 62 8.3 0.13 2.71E+21 30 -38 235 18 2.4 0.035 7.29E+20 40 -18 255 5.1 0.68 0.009 1.87E+20 50 2 275 1.5 0.20 0.003 6.25E+19 60 -26 247 0.42 0.056 0.0007 1.46E+19 80 -87 186 0.03 0.0040 0.00007 1.46E+18

Break the students into groups and then assign some of the temperature and pressure conversions to each group. Fill in the table with student values. Tell the students that--Gay-Lussac was a scientist who was interested in the properties of gases. In 1802, he set a record by ascending to 23,000 ft (~ 7 km) in a hot air balloon to investigate whether altitude affected the composition of the air. He found that it doesn’t. Because of his discovery, we can find the effective gram molecular mass (gmm) of air at all altitudes:

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gmm = 0.78(gmm N2) + 0.21(gmm O2) + 0.01(gam Ar)

gmm = 21.8 + 6.7 + 0.4 = 28.9 g/mol air We can use Avogadro’s number to find the number of gas molecules in the air at each altitude:

airofmoleculesairmolmolecules10x02.6

g9.28airmolsampleofgrams

23=

Have each group calculate the total number of molecules for the same altitudes as assigned previously. Have each group generate plots of altitude vs. temperature from atmospheric data using TI-83 or TI-82 graphing calculators. Should get something like:

Altitude vs. Temperature

0

2040

6080

100

150 200 250 300 350Temperature (K)

Alti

tude

(km

)

Discuss with the students possible reasons why the slope of the plot is not constant:

• changes as you go from one layer of the atmosphere to another • absorption of solar radiation by oxygen in middle layer results in heating of gases

Plot pressure vs. temperature from 0 to 12 km (troposphere) and pressure vs. number of molecules for the entire data set. Have the students find the best line for each plot. Discuss possible reasons why the P vs. T plot is not linear as expected (since not using a constant number of moles of gas), but P vs. number of molecules is and relate the latter result to kinetic theory of gases.

Pressure vs. Temperature

0

50

100

150

150 200 250 300 350

Temperature (K)

Pres

sure

(kPa

)

Molecules vs. Pressure

0

50

100

150

0 100 200 300

Number of molecules ( x1020)

Pres

sure

(kPa

)

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Script for Rates of Reactions PowerPoint slideshow

I’d like to talk to all of you today about dust explosions. Now, I don’t mean the proliferation indust bunnies that I see in my house, at least, when there’s a big project due. Instead, I’m talkingabout actual explosions. For example:• In 1977, a grain elevator exploded in Westwood, LA… the blast blew away the top 100 feet

of headspace off of the grain elevator, killed 36 people and injured 9. It was the worstdisaster in US history of its kind at the time.

• A little bit more recently, a flour mill in Columbus, OH, exploded killing one person andinjuring 5.

• And close to home, here in NC the West Pharmaceutical Services plant in Kinston explodedJanuary 29th, 2003. The source of the explosion was later found to be a build-up ofpolyethylene dust in the suspended ceiling.

OK…so how is it that all these different kinds of dust are exploding? Well, the key thing tothink about is that an explosion tells you that the rate of whatever reaction is occurring is very,very fast. So, we saw one example of a flour mill exploding…the flour dust in the air wasrapidly combusted. Does that mean that flour is flammable? Let’s see… [pour out small amountof flour onto Al foil and try to ignite with long handled lighter…will singe a bit, but nothingelse]. Hmm…not too exciting. Well, if you stop and think about it there’s a big differencebetween flour dust in the air and flour in a pile…Can anyone think of what that might be?[looking for surface area]. I’d like to demonstrate the flour explosion myself, but I’ve beenhaving trouble getting it to go off reliably. Fortunately, the good people in the Science Ed deptat Purdue University have been kind enough to post a video clip of a dust explosion which I’llplay for you now. [Click on icon] Say as the clip is playing: What they’re going to do is light the candle then suspend it into thelarge metal can that you see the candle hanging off of. Then they’ll take the rubber tubing that isconnected to a port a bit above the candle flame and pour some very fine, very dry lycopodiumdust (or dried moss dust) down the tubing…the dust hits the flame and voi la!

Now there aren’t that many dust explosions per year, maybe around 10 in a particularly badyear, because the dust must be present in concentrations greater than 50 g per cubic meter andthe dust particles need to have diameters less than 0.1 mm. So here’s an illustration of whathappens during a dust explosion:

• First you have the dust of the right size and concentration dispersed in a confined area• A spark or some other ignition source ignites one of the dust particles• Which then ignites all of the particles around it..• And so on…• While I couldn’t get the animation to cooperate with me on this, the rate of ignition

escalates as the ignition front advances

So, hopefully you’ve learned a little something about dust explosions, and more importantly howreaction rates are increased by an increase in surface area.

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ARTIFACTS FOR PLANNING

Meeting NCSCOS standards:In addition to meeting the specific NCSCOS standards listed in each of the lesson planspresented here (Thermochemistry: Direct Instruction, Cooperative Learning, 5-E LearningCycle), an effort has been made to address the instructional strands that are the underlying basisfor the competency goals in chemistry. Specifically, showing real world relevance (strand:Science in Personal and Social Perspectives), incorporating some of the history of chemistry intothe discussions (strand: Nature of Science), and requiring the students to predict outcomes oftheir various discovery based exercises (strand: Science as Inquiry) are examples of strategiesthat I have employed to enhance the student’s learning experience. A lesson plan entitledKinetic Theory of Gases: Cross-Curricular Integration (pp. 44-45) that incorporates bothchemical history and earth science into a lesson on gas laws has been included to illustrate myapproach.

Meeting student needs: Of the lesson plans included in this portfolio, I think the 5-E Learning Cycle lesson plan allowsthe greatest degree of flexibility for a class consisting of students with differing skill levels. Thelabels on the Ice Melter mixture and Drano are read to the students and they are asked to predictwhat they think will happen when the salts are dissolved in water. However, for any lesson planto be effective, the students must come prepared and ready to learn to class. For that reason, Ihave included a plan for instituting a token economy (pp. 46-47) in the classroom that shouldresult in the students being better prepared for class and promote improved social skills amongstthe students.

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Kinetic Theory of GasesCross-Curricular Integration

Interdisciplinary content areas: History and meteorology (Earth Science)

NC Competency Goals:Chemistry 1.06 Analyze the basic assumptions of kinetic molecular theory and its

applications.World History 7.01 Assess the degree to which discoveries, innovations, and technologies

have accelerated change. 7.02 Examine the causes and effects of scientific revolutions and cite theirmajor costs and benefits.

Earth Science 5.02 Analyze the structure of the atmosphere.5.03 Analyze weather systems.5.04 Analyze atmospheric pressure.5.05 Analyze air masses and the life cycle of weather systems.5.06 Evaluate meteorological observing, analysis, and prediction.

I. Lesson ObjectivesA. Describe the historical events that surrounded development of the gas laws.B. Identify the variables and state the relationship between them for the three gas laws. C. Connect common weather sayings to the appropriate gas law. D. Solve conceptual “real-world” problems using the gas laws.

II. Lesson PlanA. History and the Gas Laws

Evangelisto Torricelli served as Galileo’s secretary and invented the first barometerin 1643. Torricelli’s barometer used changes in the height of a column of mercuryto measure changes in air pressure.

The invention of the barometer allowed Robert Boyle to discover the relationshipbetween the pressure and the volume of a gas in 1661. Boyle’s Law states that aspressure on a gas increases its volume will decrease. Boyle also knew thattemperature affected both the pressure and volume of a gas, but couldn’t determinethe relationship since he had no way to measure temperature. It wasn’t until theinvention of the air thermometer by Guillaume Amontons in 1702 that the affect oftemperature on gas properties could be investigated.

Jacques Charles, a French physicist and chemist, released the first hydrogen filledballoon in 1783. It ascended 3 km and traveled 43 km in 2 hours. While Amontonswas actually the first to investigate the relationship between the volume andtemperature of a gas, Charles came to be credited with the discovery due to someunpublished experiments in 1787. Charles’ Law states that the volume of a gas willincrease with temperature.

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Joseph-Louis Gay-Lussac published his systematic investigations of the relationshipbetween a gases volume and temperature in 1808. However, since he referencedCharles’ unpublished results, the relationship came to be called Charles’ Law.Interestingly, Gay-Lussac was also a balloonist who had previously set a record in1802 when he ascended to 23,000 feet to investigate whether altitude affected thecomposition of the air. He found that it doesn’t. The final gas law, which states thatthe pressure of a gas will increase with temperature, is commonly attributed to Gay-Lussac; however, he is better known for his law of combining volumes.

B. Meterology and the Gas LawsIn addition to allowing Boyle to investigate gas properties, the invention of the barometer,also called a water glass, also allowed early meteorologists to predict the weather. Forexample, there’s a saying that:

“When the glass falls low, prepare for a blow;When it rises high, let all your kites fly.”

This saying tells us that you can expect a storm when there’s low air pressure, and goodweather when there’s high pressure.

A saying that predicts weather without using a barometer is:“If birds fly low, expect rain and a blow.”

The reason why this saying is a good predictor is that it’s harder for birds to fly when there’slow air pressure. Since low air pressure also accompanies storms, low flying birds are a signthat a storm is likely.

C. Practice Problems-Some real world examples

1. Why is it a bad idea to throw empty spray cans in a fire? (Hint: Are they really empty?)

2. You open a jar of instant coffee in Denver (1 mile above sea level) that was sealed at sealevel in a plant in Brazil, and it blows coffee bits up at your face. Why?

3. Arthur and Zaphod are going scuba diving in some rather cold water to try to retrieve theircopy of the Hitchhiker’s Guide to the Galaxy that has fallen overboard. While their air tankswere initially fully charged at 200 atm, Arthur notices that Zaphod’s tank reads 190 atmalmost immediately after jumping into the water. Should he drag Zaphod back up to thesurface? (Hint: The water is very cold.)

IV. SourcesA. Gas Law History excerpted from Encarta Encyclopedia and

web.fccj.org/~ethall/gaslaw/gaslaw.htm.

B. Weather sayings from wilstar.com/skywatch.htm

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Classroom Token Economy (Grades 9-12, Physical sciences)

Introduction to students. The goal behind establishing a token economy in a secondary science

education class would be to motivate students to perform tasks that are not intrinsically

motivating. Such tasks include bringing necessary materials to class everyday (paper, pencils,

textbook, calculators), completing independent homework and classwork assignments,

participating in class discussions, and remaining on task and working cooperatively when

engaged in group activities. The token economy would be introduced to the students by telling

them that I want them all to excel in my class, so I was going to give points (the tokens) to

students who completed all of these tasks on a daily basis. As each student obtains a total of 20

points, they would have to exchange their points for one of the rewards listed below. The

students can also use their tokens to purchase materials or rent a calculator for that class period

that they may have forgotten and need to complete their assignments. If there is another adult in

the room or a student assistant, their role will be to help keep track of each student’s points,

earned and exchanged, on the record sheet shown below. After two months, the efficacy of this

system will be evaluated by assessing improvement in the student’s grades on exams and

quizzes. If there has not been a significant improvement in student performance, the rules will

be modified to include points for improved test performance.

Rules

1. All students start out with 5 points. This is to avoid negative point values while the

students are becoming accustomed to the system.

2. Students receive 1 point for bringing their calculators, paper and writing utensil to class.

They do not receive a point if any one of the three items is missing.

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3. Students receive 1 point for each complete (must be neatly done or is not complete)

homework or graded classwork assignment.

4. Since cooperative effort during group work will result in a more timely completion of the

group project, all of the members of the first group that completes a group assignment

will receive 2 points each. The members of the second group to complete the group

assignment will receive 1 point each.

List of rewards

Reward Tokens required Reward Tokens required

2 sheets of paper 1 Pencil 2

Eraser 2 Candy bar 4

Rent calculator 5 Homework pass 10

1 extra credit point onnext exam 5 Drop lowest quiz

grade 20

Daily Point Tally

Monday Tuesday Wednesday Thursday FridayStudentname Prev.

M H C G M H C G M H C G M H C G M H C GEx.

A 5 1 1 1 1 1 2 1 1 1 1 1 1 1

B 5 1 1 1 1 1 1 1 1 1 1 1 2

Table Key: Prev: Point carryover from previous week, Ex: Points exchanged from reward list,M: Materials brought to class, H: completed homework, C: completed graded classwork, G:group work points

In the example above, both students started the week with 5 points. Student A was in the group

that finished first on Tuesday and received 2 points for that. Student B’s group finished second

so received 1 point for that. Student B forgot to bring their homework on Tuesday and their

materials on Wednesday. Student B exchanged some of their points for a pencil so they could

complete their classwork assignment on Wednesday.

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ARTIFACTS FOR ASSESSMENT

Formal assessment:The Free Energy Quiz from the STAD Lesson plan (pg. 32) has been included as an example ofa formal formative assessment on thermochemistry.The Enthalpy of Solution Worksheet (pp. 49-52) has also been included as an example of aformative assessment that could be used as either a formal or informal assessment of studentunderstanding depending upon the degree of difficulty that the students are having with theconcepts presented in class.

Two examples of summative traditional formal assessments have been included in this portfolio.The first example (pp. 53-55) was written to be a unit test on the kinetic theory of gases. Inaddition to multiple choice, matching and completion questions, this exam includes two free-response essay questions and two interdisciplinary bonus questions. The second exam (pp. 56-63) was written to be a unit test on thermochemistry and thermodynamics. As with the firstexample, it includes a number of different question formats but has been included in thisportfolio due to its implicit assessment of the students’ problem solving skills.

One example of an alternative summative formal assessment with its accompanying rubric isincluded on pp 64-65.

Informal assessment:A concept interview that was intended to identify potential student misconceptions regarding theproperties of gases is shown on pp. 66-67. Collecting such data about an upcoming topic can beused to tailor the lesson plans to better fit with the strengths and weaknesses of a particular groupof students.

The graphic organizer for a unit on the kinetic theory of gases (pp. 68-70) provides not only aquick means to assess student understanding of the relationships between the various gas laws,but can also be used by the student as a study aid.

Transparency 2 from the STAD Lesson plan (pp. 71-72) can be used to gather informalassessment data on student conceptual understanding of the factors that affect the spontaneity ofa physical or chemical process as well as indicators of enthalpic or entropic change. In order forthis assessment to be reflective of student understanding, input must be obtained from all of thestudents in the class. Students that seem to be having difficulty with the concepts should bequeried a number of times using different questioning formats in order to ensure that they will beable to learn the material.

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Table 1: Lattice energy values (kJ/mol)LiF 1030 NaF 910 KF 808

LiCl 834 NaCl 769 KCl 701

LiBr 788 NaBr 732 KBr 671

LiI 730 NaI 682 KI 632

MgCl2 2326 CaCl2 2223 SrCl2 2127

1. Look at the values in Table 1 and record any observations about the values you can makebelow:

Table 2: Ion hydration enthalpyCations ∆Hhyd (kJ/mol) Anions ∆Hhyd (kJ/mol)Li+ (aq) -520 F- (aq) -506Na+ (aq) -405 Cl- (aq) -361K+ (aq) -321 Br- (aq) -337

Mg2+ (aq) -1920 I- (aq) -296Ca2+ (aq) -1650 OH- (aq) -523Sr2+ (aq) -1480 NO3

- (aq) -335

2. Look at the values in Table 2 and record any observations you can make about the valuesbelow:

Enthalpy of Solution Worksheet

Name: Date:

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3. Calculate the ∆Hsoln for each of the following:

MgCl2 CaCl2 SrCl2

4. Circle which one of the following is likely to have the largest lattice energy. On the blank atthe end of each line, state the determining characteristic(s) (cation/anion, charge/size)

a) Li2SO4 Na2SO4 CaSO4

b) Na2Se Na2O Na2S

c) Al(OH)3 Al(NO3)3 AlCl3

d) AlF3 InPO4 CsCl

5. Compare the ∆Hsoln of the Group II fluorides (below) to the values you found for the Group IIchlorides in number 3.

MgF2: -17 kJ/mol CaF2: -51 kJ/mol SrF2: -14 kJ/mol

What other bonding interactions could possibly explain why the fluorides do not follow the samesize/charge trends observed for the chloride analogues? (Hint: remember that the salts are beingdissolved in water)

6. Draw an energy diagram to calculate the lattice energy for each of the following potassiumsalts:

KOH dissolved in water KNO3 dissolved in water55 kJ/mol heat released 29 kJ/mol heat absorbed

KOH (s)

E

KNO3 (s)

E

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Table 1: Lattice energy values (kJ/mol)LiF 1030 NaF 910 KF 808

LiCl 834 NaCl 769 KCl 701

LiBr 788 NaBr 732 KBr 671

LiI 730 NaI 682 KI 632

MgCl2 2326 CaCl2 2223 SrCl2 2127

1. Look at the values in Table 1 and record any observations about the values you can makebelow:

• The lattice energy decreases as the size of the anion or the cation increases.• The lattice energy seems to increase as the charge on the cation increases;

however, this might be due to the additional bonds that are formed betweenone divalent cation and two monovalent anions.

Table 2: Ion hydration enthalpyCations ∆Hhyd (kJ/mol) Anions ∆Hhyd (kJ/mol)Li+ (aq) -520 F- (aq) -506Na+ (aq) -405 Cl- (aq) -361K+ (aq) -321 Br- (aq) -337

Mg2+ (aq) -1920 I- (aq) -296Ca2+ (aq) -1650 OH- (aq) -523Sr2+ (aq) -1480 NO3

- (aq) -335

2. Look at the values in Table 2 and record any observations you can make about the valuesbelow:

• The heat of hydration become less exothermic as the size of the cation or theanion increases. This trend is also true for the divalent cations and thepolyatomic anions.

• The heat of hydration is significantly more exothermic for the divalent cationsthan it is for the monovalent cations.

Enthalpy of Solution Worksheet

Name: Answer Key Date:

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3. Calculate the ∆Hsoln for each of the following:

MgCl2 -316 kJ/mol CaCl2 -149 kJ/mol SrCl2 -75 kJ/mol

4. Circle which one of the following is likely to have the largest lattice energy. On the blank atthe end of each line, state the determining characteristic(s) (cation/anion, charge/size)

a) Li2SO4 Na2SO4 CaSO4 divalent cation

b) Na2Se Na2O Na2S smallest anion size

c) Al(OH)3 Al(NO3)3 AlCl3 smallest anion size

d) AlF3 InPO4 CsCl smallest anion size, trivalent

5. Compare the ∆Hsoln of the Group II fluorides (below) to the values you found for the Group IIchlorides in number 3.

MgF2: -17 kJ/mol CaF2: -51 kJ/mol SrF2: -14 kJ/mol

What other bonding interactions could possibly explain why the fluorides do not follow the samesize/charge trends observed for the chloride analogues? (Hint: remember that the salts are beingdissolved in water)

• Formation of hydrogen bonds between the water and the fluoride anion• Formation of HF

6. Draw an energy diagram to calculate the lattice energy for each of the following potassiumsalts:

KOH dissolved in water KNO3 dissolved in water55 kJ/mol heat released 29 kJ/mol heat absorbed

cation

-523 kJ/mol

-321 kJ/mol

KOH (s)

E

-55 kJ/mol∆Hsoln

K+ (aq) + OH-

K+ + OH -

K+ (aq) + OH- (aq)

U = 789 kJ/mol

KNO3 (s)

E

+29 kJ/mol

K+ + NO3-

K+ (aq) + NO3-

-321 kJ/mol

-335 kJ/molK+ (aq) + NO3

- (aq)

U = 685 kJ/mol

Page 53: Cooperative Learning in Chemistry

Sample Unit Test Kinetic Theory of Gases

I. Warm Up: Please write the letter of the best choice in the blank to the left. _____ 1. Why is the Kelvin temperature scale used when solving gas law problems?

a) to make the math more complicated b) it’s used for historical reasons c) to prevent negative volumes d) to prevent positive pressures

_____ 2. How would the number of molecules of two different gases compare if their partial

pressures in the same container were identical? a) They would be the same b) There would be fewer molecules of the heavier gas c) There would be more molecules of the heavier gas d) There’s not enough information to tell

_____ 3. What happens to gas molecules when the gas is compressed?

a) The molecules flatten out b) The molecules get closer together c) The molecules slow down d) The molecules stop moving

_____ 4. Boyle’s Law states that the pressure and volume of a gas is:

a) conversely proportional b) directly proportional c) inversely proportional d) reversely proportional

_____ 5. Avogadro’s hypothesis states that one mole of gas at STP will have a volume of: a) 2.24 mL b) 22.4 mL c) 2.24 L d) 22.4 L

II. Word Play: Match each definition below to the word bank at the right. _____ 1. Force per unit area a) Charles’ Law _____ 2. Movement of gases from regions of high concentration

to lower concentration b) Kinetic Theory of

Gases _____ 3. Volume and temperature of a gas are directly related c) Ideal Gas Law _____ 4. As the temperature of a gas increases, the average

speed of the gas molecules increases d) Diffusion

_____ 5. PV=nRT e) Pressure

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III. The Crunch: Please show your work in the space provided 1. What is the volume of 0.25 moles of a gas at STP? 2. If the atmosphere is 78% N2, 21% O2 and 1% Ar, what is the partial pressure of each gas at

sea level (101.3 kPa)? 3. A balloonist puts 1.25 x 105 L of air into her balloon at 25 °C at atmospheric pressure. The

air is heated to 200 °C at constant pressure. What is the volume of air at this temperature? 4. Which gas effuses faster-hydrogen or chlorine? How much faster? 5. A 45 g sample of a volatile liquid is completely vaporized in a 10 L container at 100 °C and

100 kPa. What is the molecular mass of the gas?

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IV. Home stretch: Please answer the following conceptual questions as completely as possible.

1. What are three differences between ideal gases and real gases? 2. Draw a picture of a balloon of gas at STP at the molecular level. Draw another picture that illustrates how conditions change at the molecular level when the temperature of the gas is decreased, and explain what happens.

V. Bonus questions:

1. Why didn’t Robert Boyle investigate the effect of temperature on the volume of a gas?

2. You’re out at Jordan Lake and you notice a flock of ducks flying across the lake barely above the surface. What is likely to happen?

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1

THERMOCHEMISTRY: EXAM

Part I. Warm-up exercises (Fill in the blank and multiple choice)

1. Label each of the following as either a physical change (P) or a chemical change (C) (4 pts):

_____ CO2 (s) → CO2 (g) _____ C6H12O6 (s) + O2 (g) → CO2 (g) + H2O (g)

_____ 2H2 (g) + O2 (g) → 2H2O (g) _____ NaCl (s) → NaCl (l)

2. Indicate whether each of the following physical and chemical changes results in a positive(+), negative (-) or negligible (0/+ or 0/-) ∆S value (4 pts):

_____ CO2 (s) → CO2 (g) _____ C6H12O6 (s) + O2 (g) → CO2 (g) + H2O (g)

_____ 2H2 (g) + O2 (g) → 2H2O (g) _____ NaCl (s) → NaCl (l)

3. Circle which one of the following is likely to have the largest lattice energy. List thedetermining characteristics (cation/anion, charge/size) in the accompanying blank (5 pts):

a) Li2SO4 Na2SO4 CaSO4

b) Al(OH)3 Al(NO3)3 AlF3

c) AlF3 InPO4 CsCl

_____4. An endothermic process (2 pts):

a) absorbs energy from the surroundings b) releases heat to the surroundings

c) has a negative ∆H d) both (a) and (c)

_____5. Spontaneous reactions (2 pts):

a) Always release heat energy b) Always have a negative ∆H

c) Always have a negative ∆G d) Always have a positive ∆H

_____6. The mathematical expression that relates free energy, enthalpy and entropy is (2 pts):

a) ∆H = ∆G – T∆S b) ∆S = ∆H – T∆G c) ∆G = ∆H – T∆S d) ∆H = ∆S – T∆G

_____7. Sodium acetate dissolves readily in water according to the following equation:

NaC2H3O2 (s) → NaC2H3O2 (aq) ∆H = -17.3 kJ/mol

The temperature of the water used to dissolve the sodium acetate will (2 pts):

a) not change b) increase

c) decrease d) impossible to tell from the information given

Name: Date:

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2

8. For each of the following, indicate the sign of ∆G and ∆H (5 pts).

∆G ∆H

a) Dissolving a solid in water results in a ∆T = +3°………….

b) The melting of ice at room temperature…………………...

c) Condensation of water vapor at 120°C……………………

d) Condensation of water vapor at 80°C……………………..

e) Sublimation of dry ice at room temperature………………

Part II. Time to turn up the heat! (Calculations-Please show your work in the space provided.)

Questions 1-4 refer to flatulite (H2PEw, gfm = 61.90 g/mol), a fictitious flammable compoundfrom the planet Flatulon. It has a melting point of 132°C and a boiling point of 215°C.

1. A 135-g piece of solid flatulite is heated to 90.5 °C and then dropped into 75.5 g of water thatis initially at 21.3 °C. What is the specific heat of solid flatulite if the final temperature ofthe water is 32.3 °C? (CpH2O(l) = 4.18 J/g°C) (4 pts)

2. What is the molar heat of fusion of flatulite if it requires 311 J of heat energy to completelymelt 55.7 g of the solid? (2 pts)

3. What is the molar heat of condensation of flatulite if it requires 658 J of heat energy tocompletely vaporize 112 g of liquid flatulite? (3 pts)

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3

4. The combustion of flatulite is an exothermic reaction.

H2PEw (s) +3O2 (g) → 2H20 (l) + P2O5 (s) + 2Ew2O (s) ∆Hrxn = -139 kJ

a) Calculate the amount of heat liberated when 4.79 g of H2PEw reacts with excess oxygen.(3pts)

b) What would the ∆Hrxn be if gaseous water were formed instead of liquid water? [∆Hf

(H2O(l) = -285.8 kJ/mol), ∆Hf (H2O(g) = -241.8 kJ/mol] (3 pts)

5. Find the standard enthalpy of formation of phosphorus pentachloride from its elements (5pts).

2P (s) + 5 Cl2 (g) → 2PCl5 (s)Use the following thermochemical equations.

PCl5 (s) → PCl3 (g) + Cl2 (g) ∆H = 87.9 kJ

2P (s) + 3Cl2 (g) → 2PCl3 (g) ∆H –574 kJ

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4

III. The final cool-down (diagramming)

1. Draw an energy diagram that could be used to calculate the lattice energy for each of thefollowing potassium salts. Indicate on the diagram the additional pieces of informationneeded to complete the calculation. (10 pts)

KOH dissolved in water55 kJ/mol heat released

2. Draw a diagram that illustrates the direction of heat flow, and the temperature of the systemvs. the surroundings for an endothermic reaction. Please identify the system and thesurroundings on the diagram. (4 pts)

KOH (s)

E

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THERMOCHEMISTRY: EXAM(60 pts total)Part II. Warm-up exercises (Fill in the blank and multiple choice)

1. Label each of the following as either a physical change (P) or a chemical change (C) (4 pts):

P CO2 (s) → CO2 (g) __C C6H12O6 (s) + O2 (g) → CO2 (g) + H2O (g)

C 2H2 (g) + O2 (g) → 2H2O (g) P NaCl (s) → NaCl (l)

2. Indicate whether each of the following physical and chemical changes results in a positive(+), negative (-) or negligible (0/+ or 0/-) ∆S value (4 pts):

+ CO2 (s) → CO2 (g) + H12O6 (s) + O2 (g) → CO2 (g) + H2O (g)

2H2 (g) + O2 (g) → 2H2O (g) 0/+ NaCl (s) → NaCl (l)

3. Circle which one of the following is likely to have the largest lattice energy. List thedetermining characteristics (cation/anion, charge/size) in the accompanying blank (5 pts):

a) Li2SO4 Na2SO4 CaSO4 cation charge

b) Al(OH)3 Al(NO3)3 AlF3 anion size

c) AlF3 InPO4 CsCl cation charge, anion size

_____4. An endothermic process (2 pts):

b) absorbs energy from the surroundings b) releases heat to the surroundings

d) has a negative ∆H d) both (a) and (c)

_____5. Spontaneous reactions (2 pts):

c) Always release heat energy b) Always have a negative ∆H

c) Always have a negative ∆G d) Always have a positive ∆H

_____6. The mathematical expression that relates free energy, enthalpy and entropy is (2 pts):

a) ∆H = ∆G – T∆S b) ∆S = ∆H – T∆G c) ∆G = ∆H – T∆S d) ∆H = ∆S – T∆G

_____7. Sodium acetate dissolves readily in water according to the following equation:

NaC2H3O2 (s) → NaC2H3O2 (aq) ∆H = -17.3 kJ/mol

The temperature of the water used to dissolve the sodium acetate will (2 pts):

a) not change b) increase

c) decrease d) impossible to tell from the information given

Name: Answer Key Date:

A

C

C

B

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6

8. For each of the following, indicate the sign of ∆G and ∆H (5 pts).

∆G ∆H

a) Dissolving a solid in water results in a ∆T = +3°…………. Negative Negative

b) The melting of ice at room temperature…………………... Negative Positive

c) Condensation of water vapor at 120°C…………………… Positive Negative

d) Condensation of water vapor at 80°C…………………….. Negative Negative

e) Sublimation of dry ice at room temperature……………… Negative Positive

Part II. Time to turn up the heat! (Calculations-Please show your work in the space provided.)

Questions 1-4 refer to flatulite (H2PEw, gfm = 61.90 g/mol), a fictitious flammable compoundfrom the planet Flatulon. It has a melting point of 132°C and a boiling point of 215°C.

1. A 135-g piece of solid flatulite is heated to 90.5 °C and then dropped into 75.5 g of water thatis initially at 21.3 °C. What is the specific heat of solid flatulite if the final temperature ofthe water is 32.3 °C? (CpH2O(l) = 4.18 J/g°C) (4 pts)

qw= (4.18 J/g°C)(75.5 g)(32.3°C – 21.3°C) = 3471 J (1 pt)qFt = -qw = -3471 J (1 pt)Cp(Ft) = -3471 J/[(32.3°C – 90.5°C)(135 g) = 0.442 J/g°C (2 pts)

2. What is the molar heat of fusion of flatulite if it requires 311 J of heat energy to completelymelt 55.7 g of the solid? (2 pts)

55.7 g/(61.9 g/mol) = 0.8998 mol Ft (1 pt)

∆Hfus = 311 J/0.900 mol = 346 J/mol (1 pt)

3. What is the molar heat of condensation of flatulite if it requires 658 J of heat energy tocompletely vaporize 112 g of liquid flatulite? (3 pts)

112 g/61.9 g/mol = 1.809 mol Ft (1 pt)∆Hconds = -∆Hvap (1 pt for keeping track of sign change)∆Hconds = -658J/1.809 mol = -364 J/mol (1 pt)

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4. The combustion of flatulite is an exothermic reaction.

H2PEw (s) +3O2 (g) → 2H20 (l) + P2O5 (s) + 2Ew2O (s) ∆Hrxn = -139 kJ

a) Calculate the amount of heat liberated when 4.79 g of H2PEw reacts with excess oxygen.(3pts)

4.79 g/61.9 g/mol = 0.0774 mol Ft (1 pt)q = (0.0774 mol Ft)(-139 kJ/mol) = -10.8 kJ (1 pt)10.8 kJ of heat are liberated (1 pt for equating neg sign with heat release)

d) What would the ∆Hrxn be if gaseous water were formed instead of liquid water? [∆Hf

(H2O(l) = -285.8 kJ/mol), ∆Hf (H2O(g) = -241.8 kJ/mol] (3 pts)

∆Hrxn = Σ∆H(products) - Σ∆H(reactants)-139 kJ = [2(Hf(H2Ol) + Hf(P2O5) + 2Hf(Ew2O)]-[Hf(H2PEw)] (1 pt)Subtract off contribution from 2 equivalents of liquid water, add in contributionfrom 2 equivalents of gaseous water (1 pt for including coefficients)

New ∆Hrxn = -95 kJ (1 pt)

5. Find the standard enthalpy of formation of phosphorus pentachloride from its elements (5pts).

2P (s) + 5 Cl2 (g) → 2PCl5 (s)Use the following thermochemical equations.

PCl5 (s) → PCl3 (g) + Cl2 (g) ∆H = 87.9 kJ

2P (s) + 3Cl2 (g) → 2PCl3 (g) ∆H –574 kJ

2x [Cl2 (g) + 2PCl3 (g) → PCl5 (s)] 2x (-87.9 kJ) (1 pt for neg sign, 1 pt for coeff)2P (s) + 3Cl2 (g) → 2PCl3 (g) -574 kJ (1 pt)

2P (s) + 5 Cl2 (g) → 2PCl5 (s) -749.8 kJ = ∆Hrxn (1 pt)

∆Hf(PCl5) = ½ ∆Hrxn = -374.9 kJ (1 pt)

+

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8

III. The final cool-down (diagramming)

1. Draw an energy diagram that could be used to calculate the lattice energy for potassiumhydroxide. Identify the species present at each energy level, and label the enthalpy changes.Indicate on the diagram the additional pieces of information needed to complete thecalculation. (11 pts)

KOH dissolved in water55 kJ/mol heat released

2. Draw a diagram that illustrates the direction of heat flow, and the temperature of the systemvs. the surroundings for an endothermic reaction. Please identify the system and thesurroundings on the diagram. (3 pts)

KOH (s)

E

-55 kJ/mol∆Hsoln

K+ (aq) + OH-

K+(g) + OH–(g)

K+ (aq) + OH- (aq)

U ∆Hhyd (K+)

∆Hhyd (OH-)

1 pt for each arrow in correctdirection (4 pts)

1 pt each for labeling horizontallines (3 pts)

1 pt for labeling each enthalpychange (4 pts)

system

surroundings

q

Tsys < Tsurr

1 pt for q arrow direction1 pt for sys/surr ID1 pt for correct T relationship

Page 64: Cooperative Learning in Chemistry

THERMOCHEMISTRYAlternative Assessment

Part I. Turning up the heat! (40 points) We have been studying how heat and chemical changeare related. Use your knowledge of enthalpy, specific heat and calorimetry to:

a) Design an experiment to measure the heat of combustion of potato chips.b) Assuming 100% combustion, calculate the magnitude of the experimental values that you

will need to collect.c) Predict potential sources of error that would lead to less than 100% combustion.d) Describe in general terms one other technique that could be used to examine the efficiency of

potato chip combustion.

RubricPoints → 6 8 10

Experimentaldesign

• The materials andequipment list is missing.

• The experimental apparatusis not diagrammed.

• The experimental procedureis incomplete.

• Most of the materials andequipment needed are listed

• The experimental apparatusis missing some of the finerdetails

• The experimental procedureis complete, but notsufficiently detailed foranother person to performthe experiment.

• All of the materials andequipment needed is listed• The experimental apparatus

is clearly diagramed• The experimental procedure

provides sufficient detail forsomeone else to perform theexperiment

Calculations • While the ingredient listfrom a package of potatochips has been included,there were no calculationsdone that used it as aresource

• Little if no computationalrationale is provided tosupport the quantity of thematerials to be used in theexperiment

• The ingredient list from apackage of potato chips hasbeen used as a resource

• An attempt has been madeto address the effect that thefat and salt content willhave on the results, but noactual calculations wereperformed

• Most of the quantities usedfor the experiment willyield experimentallyrealizable values

• The ingredient list from apackage of potato chips hasbeen used as a resource• The effect of the salt and fat

content of the chips has beenconsidered• The quantities of materials

used for the experiment havebeen adjusted to achieveexperimentally realizablevalues

Sources oferror

• List is missing, or includesonly human error

• List is comprehensive, butlacks a prediction ofpotential impact upon finalresults

• List is comprehensive andincludes analysis of how thepotential error would impactthe final results

Follow-upexperiment

• A duplicate of the firstexperiment with only thequantities changed

• Missing essential detailneeded to evaluatefeasibility of approach

• Sufficiently detailed toevaluate the feasibility of theapproach

Name: Date:

Page 65: Cooperative Learning in Chemistry

Part II. Chilling out. (20 points) Keep a journal for three days (5 points for journal) and recordany examples of endothermic physical or chemical processes that you encounter during that time.Be sure to include estimates of the duration of the observation, temperature changes, and anyother observations that might enhance your data analysis. After collecting your data, compileyour observations into physical and chemical processes then rank them within their group fromthe greatest amount of energy transferred to the least. Discuss any trends that you can deducefrom your two lists.

RubricPoints → 3 4 5Journal • Observations do not cover

all three days, and lackdetail

• Observations have beenlogged for the three days,but are missing most if notall of the additional datathat needed to be collected

• Detailed observations havebeen logged for all threedays

DataOrganization

• The data has not beenorganized into any coherentstructure

• Either the division of theobservations between thetwo main categories isincorrect

OR

• The ranking of the extent ofenergy transfer is incorrect.

• The chemical and physicalprocesses have beencorrectly separated into twolists. • The observations have been

appropriately rankedaccording to energytransferred within each maincategory

Trends • Little or no effort was madeto analyze the data

• Only information presentedin class was used to analyzethe data

• Outside information hasbeen utilized in order toanalyze the data

Page 66: Cooperative Learning in Chemistry

The concept interview subjects were four female Apex High Academic Chemistry students. They had finished their lab earlier thanthe rest of the class and were available for the interview. The class has not yet studied the properties of gases. The interview was cut alittle short because there was only 20 minutes left before class ended.

Interview

Blow up a balloon, hold it closed and then ask,

Student# 1 2 3 4

Can you draw a picture ofwhat you think it looks likeinside? (They drew:)

A balloon with dots inthe center

A balloon with dotsand shading

A balloon with dotsevenly distributedthroughout

A balloon with dotsevenly distributedthroughout

What would it sound like? Nothing A hum Nothing Like waves

Blow up the balloon some more and ask,

What would your picture looklike now? (They drew: )

A larger balloon withmore dots, but stillmostly in center

A larger balloon withdarker dots andshading

A larger balloon withdots evenly distributedthroughout

A larger balloon withdots evenly distributedthroughout

What would happen if I keptblowing up the balloon? It would pop It would pop It would pop It would pop

Why would it do that?Because the outsidewould be stretched toothinly

I don’t know I don’t know Because there’s toomuch stuff inside

Page 67: Cooperative Learning in Chemistry

Let’s use something a little stronger than a balloon. Let’s imagine that I’ve got a very heavy duty air pump attached to a steel can andthat I pump it up to about 30 psi, which is about the same as an automobile tire.

What would happen if Iwere to heat it up a littlebit?

Nothing Nothing Nothing The stuff inside wouldmove around faster

What would happen if Iwere to heat it up a lot? The can would melt Nothing Nothing

The movement inside ofthe can would get so fastthat the can would burst.

To student#4: Whywould the can burst?

The molecules inside aremoving around so much,and they’re trying to stayaway from each other atthe same time.

Now imagine that I have a hula hoop with a rubber sheet stretched over it.

What would happen ifwe were all to throwtennis balls at the rubbersheet?

The center would stretchout.

The center would stretchout.

The center would stretchout.

The center would stretchout.

What would happen ifwe were to throw theballs harder?

It would stretch evenmore Stretch more Stretch more Stretch more

Page 68: Cooperative Learning in Chemistry

Grade level: High school chemistryConcept: Properties of gasesLesson objective:After completing this lesson, the students will be able to

• Compile a list of the gas laws containing both the formula and the associatedname

• Complete the flow chart using their compiled list• Use the flow chart to solve word problems involving the gas laws

Lesson Plan:After being introduced to the gas laws (Boyle’s, Charles’, Gay-Lussac, combined andideal), the students will form groups of 2-3 people and compile a list of the named gaslaws and their formulae. After having the list checked for accuracy, the group will begiven a copy of the graphic organizer (flow chart). The group will then utilizediscussion and their list to complete the flow chart. Completion of the flow chart willfocus the students’ attention on the interrelationships of the four variables (P, V, Tand n) encountered in gas law word problems. The students will then check theirown flow chart for accuracy by using it to solve a series of gas law word problems.After all of the groups have completed the word problems, a student from each groupwill be asked to explain how they went about solving one of the problems. Groupsthat had difficulty completing either the flow chart or solving the problems will beasked to explain what they were able to accomplish and then provided with sufficientcoaching from either the teacher or from other groups (who have already presented aproblem and upon the teacher’s request to provide assistance) to finish workingthrough the problem.

Page 69: Cooperative Learning in Chemistry

Pressure? No

# moles?

Yes

No

Yes

No

2

2

1

1TP

TP

=

No

Yes Yes

Does the probleminvolve:

(Student version)

Page 70: Cooperative Learning in Chemistry

Pressure? Temperature? 2211 VPVP =No

# moles?2

22

1

11TVP

TVP

=

Yes

No

nRTPV =

Yes

2

2

1

1TV

TV

=

No

Volume?

2

2

1

1TP

TP

=

No

Yes Yes

Does the probleminvolve:

(Completed version)

Charles’ Law Gay-Lussac’s Law Combined Gas Law

Boyle’s Law

Ideal Gas Law

Page 71: Cooperative Learning in Chemistry

∆G ∆H ∆S

H2O (s) → H2O (l)T = 293 K

H2O (s) → H2O (l)T = 253 K

H2O (l) → H2O (g)T = 383 K

NaOH (s) → Na+ (aq) + OH- (aq)∆T > 0

CsCl (s) → Cs+ (aq) + Cl- (aq)∆T < 0

2C2H2 (g) + 5O2 (g) → 4CO2 (g) + 2H2O (g)∆T > 0

Transparency 2

Page 72: Cooperative Learning in Chemistry

∆G ∆H ∆S

H2O (s) → H2O (l)T = 293 K − + 0/+

H2O (s) → H2O (l)T = 253 K + + 0/+

H2O (l) → H2O (g)T = 383 K − + +

NaOH (s) → Na+ (aq) + OH- (aq)∆T > 0 − − +

CsCl (s) → Cs+ (aq) + Cl- (aq)∆T < 0 − + 0/+

2C2H2 (g) + 5O2 (g) → 4CO2 (g) + 2H2O (g)∆T > 0 − − −

Key: Transparency 2