Name ...



Name: _______________________________________________ Period: _________

Thermochemistry

Conservation of Energy

Pre-AP

EVERY DAY you need to bring your calculator!!

| | |In Class |Homework (you write in the assignment for the |

| | | |night) |

|Monday |3/22 |Intro to energy & three types of heat; Endothermic/exothermic; Reaction | |

| | |coordinates | |

|Tuesday |3/23 |Law of conservation of energy; Heat (q) of substance; Specific heat (C); | |

| | |Heat capacity (HC) | |

|Wednesday |3/24 |Identify an unknown metal LAB | |

|Thursday |3/25 |Enthalpy (H); Heat (q) of reaction | |

|(late start) | | | |

|Friday |3/26 |Heat (q) of combustion LAB | |

|Monday |3/29 |Heat (q) of calorimeter; post-lab discussion | |

|Tuesday |3/30 |QUIZ (specific heat, heat of reaction) | |

| | |Heat (q) of solution | |

|Wednesday |3/31 |Heat of solution & reaction LAB; | |

| | |Phase change LAB | |

|Thursday |4/1 |Heat (q) of phase change | |

|Friday |4/2 |Holidaaay! Celebraaate! | |

|Monday |4/5 |Hess’ law | |

|Tuesday |4/6 |Enthalpy of formation; Review | |

|Wednesday |4/7 |New unit: Solutions | |

|Thursday |4/8 |New unit: Solutions | |

|Friday |4/9 |TEST | |

*Fill in the schedule and assignments every day in class as we progress. Be flexible!

Summary table: Types of heat (q). m=mass, C=specific heat capacity; HC=heat capacity, H=enthalpy

|qsubstance |qprocess |qcalorimeter |

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warm ups warm ups

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Energy

Vocab: energy, calorie, law of conservation of energy (LoCE), heat, q, endothermic, exothermic, reaction coordinate diagram

Objectives (students will be able to):

➢ define vocab terms

➢ list types of energy

➢ apply the law of conservation of energy to a hamburger/veggie burger being eaten

➢ explain the key points of a reaction coordinate diagram

➢ indicate if a reaction is endothermic or exothermic using the reaction coordinate diagram

Notes/Problem Workspace:

[pic]

|Substance |Specific Heat Capacity |

| | |

| |at 25oC in J/goC |

|H2 gas |14.267 |

|He gas |5.300 |

|H2O(l) |4.184 |

|lithium |3.56 |

|ethyl alcohol |2.460 |

|ethylene glycol |2.200 |

|ice @ 0oC |2.010 |

|steam @ 100oC |2.010 |

|vegetable oil |2.000 |

|sodium |1.23 |

|air |1.020 |

|magnesium |1.020 |

|aluminum |0.900 |

|Concrete |0.880 |

|glass |0.840 |

|potassium |0.75 |

|sulfur |0.73 |

|calcium |0.650 |

|iron |0.444 |

|nickel |0.440 |

|zinc |0.39 |

|copper |0.385 |

|brass |0.380 |

|sand |0.290 |

|silver |0.240 |

|tin |0.21 |

|lead |0.160 |

|mercury |0.14 |

|gold |0.129 |

q and Specific Heat

Vocab: heat, heat capacity, specific heat, q, law of conservation of energy (LoCE), thermal equilibrium

Objectives (students will be able to):

➢ define vocab terms

➢ calculate q for a substance changing temperature

➢ use the LoCE to set up equations and calculate an unknown variable (SH/temp/mass)

Notes/Problem Workspace:

Identify an unknown metal LAB:

Mass of water ___________

Ti of water in the calorimeter ___________

Ti of metal ___________

Tf of metal and water ___________

Mass of metal ___________

Calculate specific heat of metal:

How would the calculated value be different if it took too long to transfer the metal to the calorimeter?

Problems for Heat Energy and Law of Conservation of Energy:

1. When will q be positive? When will it be negative? What information in a question will give you clues about q being positive or negative?

2. Calculate the energy required to heat a 10-gram piece of gold from 25°C to 38°C. (answer = 16.77 J)

3. Calculate the energy (in J and kJ) required to cool a piece of glass from 100°C to 0°C. The glass weighs 1.34 kg. (answer = -112560 J or -112.560 kJ)

4. What is thermal equilibrium?

5. A small piece of unknown metal (0.5 g) is placed in a 1.7-kg beaker of olive oil (SHC = 1.97 J/g°C) at room temperature (25°C). The metal cools from 88.00°C to 25.03°C. What is the specific heat capacity of the unknown metal? (answer = 3.19 J/gºC)

6. A piece of iron at -50°C is heated by placing the 2-g metal chunk in a large beaker of water at room temperature (25°C). The final temperature is 24.8°C. What volume of water (in mL) was in the beaker? (answer = 79 mL H2O)

Enthalpies of Reactions and Heats of Reactions

Vocab: enthalpy, enthalpy of reaction (ΔHrxn), enthalpy of combustion, enthalpy of neutralization, enthalpy of formation, heat of reaction (qrxn), calorimetry

Objectives (students will be able to):

➢ define vocab terms

➢ identify a reaction as endothermic or exothermic based on its enthalpy of reaction

➢ use the LoCE to set up equations and calculate an unknown variable (SHC/temp/mass/enthalpy)

➢ use stoichiometry and enthalpies of reactions to calculate an unknown (energy/mass/mole)

Notes/Problem Workspace:

Nitroglycerine, C3H5(NO3)3, is the major component in dynamite. It explodes according to the reaction below with ∆H°rxn= -5692 kJ/mol rxn.

4 C3H5(NO3)3(l) ( 6N2(g) + O2(g) + 10H2O(g) + 12CO2(g)

a. How many kilojoules are released when 125 grams of nitroglycerine react?

b. How many kilograms of nitroglycerine explode react to release 1.00 MJ?

c. If 125 grams of nitroglycerine react to heat water in a can initially at 25°C to 29°C, what was the initial volume of water in the can?

Problems for Stoichiometry using Enthalpies of Reactions:

7. Propane, C3H8, is a common fuel, which we often use for portable stoves and lanterns. When propane burns in air, the ∆H°rxn= -2220 kJ/mol rxn.

a. Write the balanced reaction for the combustion of propane.

b. When this reaction provides 375 kJ of heat, how many grams of CO2 form? (answer = 22.3 g CO2)

c. How many kilojoules of heat are released when 375.0 grams of propane burn? (answer = -18920 kJ)

d. If this reaction is used to heat 87 mL of water initially at 0°C in a can, what would the final temperature of the can be using the information provided in part b? Would I want to drink the water? (answer = 1031ºC)

8. When 9.50 grams of nitrogen react with oxygen to form nitrogen dioxide, the system absorbs 11.7 kilojoules from the surroundings.

a. Write the balanced chemical equation. Is this reaction endothermic or exothermic? How do you know?

b. What is the ∆H°rxn for this process in kJ/mol NO2? (answer = 17.2)

9. The heat of formation of copper(I) chloride is –137 kJ/mol rxn.

a. Write the balanced chemical equation. [Don’t forget the definition of a heat of formation!!!]

b. How many joules are released when 4.46 grams of copper react with excess chlorine to produce copper(I) chloride? (answer = -9615 J)

c. If this energy was used to heat a 40-gram piece of glass at 21°C, what would the final temperature of the glass be? (answer = 307ºC)

10. Hydrogen sulfide reacts with oxygen to produce water and sulfur dioxide with ∆H°rxn= -3330 kJ/mol rxn

a. Write the balanced chemical reaction.

b. What is the heat change in kJ if 1.31 moles of hydrogen sulfide are reacted? (answer = -2181)

c. How man grams of water are formed when the reaction releases 270 kJ? (answer = 2.92 g water)

d. If I am trying to heat a 150-g piece of zinc from 30°C to 80°C, how much hydrogen sulfide would I need (reacted with excess oxygen)? (answer = 0.06 g hydrogen sulfide)

11. A jumbo Hill Country Fare brand marshmallow weighing 7 grams is burned. The flames heat a can holding 200 mL of water initially at 10°C. Assume only 35% of the marshmallow mass combusts. When the marshmallow stops burning, the water reaches a peak temperature of 49°C. Calculate the amount of food calories per gram of marshmallow released from the combustion reaction. [Remember that a food calorie is 1000 calories.] (answer = 3.2 food cal/gram)

FOOD CALORIMETRY

I. OBJECTIVE

II. CHEMICALS/REAGENTS: none

III. DRAWING AND APPARATUS

IV. PROCEDURE (paraphrase the steps below)

1. Collect all materials and bring them to your lab station. (Clean the inside of your coke can if needed. It’s okay if the outside is blackened.)

2. Weigh the clean, dry can and record its mass in your data table. Add 200mL of water (measured using a graduated cylinder) to the soda can. Weigh and record the mass in the data table. Hang the can on the ring stand by placing the colored pencil through the coke can tab. This apparatus is your calorimeter.

3. Straighten part of the paper clip and leave one end bent. Turn the cork upside down and poke the straight end into the cork. The bent or looped end of the clip (it should be flat) is the platform for the food.

4. Use a thermometer clamp to suspend the thermometer in the can. It should be in the water but not touching the sides or bottom of the can. Before going any further, check the apparatus to make sure that everything is secured.

5. Choose a piece of food, measure its initial mass and record in the data table.

6. Take an initial temperature reading of the water in the can and record in the data table.

7. Place the food on the paperclip stand. Use a match to set the food on fire. This may take several tries. Closely observe the food as it burns. If the fire is not on the bottom of the can, move the cork so that the flame is heating the can. If the food falls off the stand, replace it on the paperclip stand and re-light immediately. If it takes too long the water in the can may start to cool. If this happens, re-start your experiment from step 5.

8. As soon as the food stops burning, observe the temperature. When the temperature hits its peak temperature, record the temperature in your data table (you will know it’s the peak temperature because the temperature will start to decrease again).

9. As soon as the food has cooled, carefully take the food to the scale. Place a paper towel on the scale and tare the scale. Place the burned remains onto the paper towel, and record the final food mass in the data table. Throw the remains and the paper towel in the trashcan.

10. Follow your teacher’s instructions for the number of trials you are to perform. In order to ensure consistency between trials, you may need to get new water, wash your can, or cool your can by running cool water over it. When you are done with all your trials, clean up your area thoroughly and return all items to their original location.

V. EQUATIONS: omit (combustion occurs but the formula for the cheeto is unknown)

VI. DATA TABLES: create your own based on procedure

VII. QUESTIONS AND CALCULATIONS

1. Calculate change in water temperature, Δt, for each sample.

2. Calculate the mass (in g) of the water heated for each sample.

3. Use the results of Steps 1 and 2 to determine the heat energy gained by the water (in J).

4. Calculate the mass (in g) of each food sample actually combusted.

5. Use the results of Steps 3 and 4 to calculate the energy content (in J/g) of each food sample.

6. Convert the energy content to kilocalories per gram.

7. Using the nutrition information on the food package, calculate the percent error.

VIII. CONCLUSIONS: thoroughly answer the following:

Discuss the sources of heat loss during this experiment. Suggest a better design to eliminate error.

Calorimeters

Vocab: calorimeter, bomb calorimeter, qcal

Objectives (students will be able to):

➢ define vocab terms

➢ describe the parts of a calorimeter and explain how it works

➢ use the LoCE to set up equations and calculate an unknown variable (SHC/temp/mass/enthalpy)

Notes/Problem Workspace:

A 2.2-g sample of quinone, C6H4O2, is burned in a bomb calorimeter whose total heat capacity is 7.854 kJ/°C. The temperature of the calorimeter increases from 23.44°C to 30.57°C. What is the enthalpy of combustion per mole reaction?

Problems for Calorimetry using Enthalpies of Reactions (with heats of reaction):

12. A 1.8 g sample of octane, C8H18, was burned in a bomb calorimeter whose total heat capacity is 11.66 kJ/°C. The temperature of the calorimeter plus contents increased from 21.36°C to 28.78 °C. What is the enthalpy of combustion per mole reaction? (answer = -5479 kJ/mol rxn)

13. The heat of combustion of glucose, C6H12O6, is -15.57 kJ/g. A 2.5-g sample of glucose is burned in a bomb calorimeter. The temperature of the calorimeter increased from 20.55°C to 23.25°C.

a. What is the total heat capacity of the calorimeter? (answer = 14.4 kJ/ºC)

b. If the calorimeter contained 2.7 kg of water, what is the heat capacity of the dry calorimeter? (answer = 3.06 kJ/ºC)

c. What temperature increase would be expected in this calorimeter if the glucose sample had been combusted when the calorimeter contained 2 kg of water?

14. The heat of neutralization for sulfuric acid reacting with potassium hydroxide is -221.6 J/mol rxn.

a. Why is this called a neutralization reaction?

b. Write the balanced chemical reaction.

c. A reaction takes place between 20 mL of 2M sulfuric acid and 5 grams of potassium hydroxide. How much energy (in J) is released? (answer = -8.864 J)

d. An unknown mass of water is heated with the energy released in part c. The temperature rises by 0.02°C. Calculate the mass of water. (answer = 106 mL of water)

Enthalpies of Solution and Heats of Solution

Vocab: solution, physical change, heat of solution, enthalpy of solution, dissociation, solvation, solute, solvent, molarity

Objectives (students will be able to):

➢ define vocab terms

➢ use the LoCE to set up equations and calculate an unknown variable (SHC/temp/mass/enthalpy)

Notes/Problem Workspace:

| |Enthalpies |

|ammonium nitrate |+6.14 kcal/mol |

|ammonia |-7.29 kcal/mol |

|potassium hydroxide |-13.77 kcal/mol |

|cesium hydroxide |-17.10 kcal/mol |

|sodium chloride |+0.93 kcal/mol |

|potassium chlorate |+9.89 kcal/mol |

|acetic acid |-0.360 kcal/mol |

|hydrochloric acid |-17.89 kcal/mol |

|Barium chloride (anhydrous) |-8.67 kJ/mol |

|Barium chloride dihydrate |+20.58 kJ/mol |

|Magnesium sulfate heptahydrate |+16.11 kJ/mol |

|Ammonium nitrate |+26.44 kJ/mol |

|Na2S2O3 · 5H2O |+47.40 kJ/mol |

When a 6.50 g sample of solid sodium hydroxide dissolves in 100 g of water in a coffee-cup calorimeter, the temperature rises from 21.6°C to 37.8°C. Calculate the change in enthalpy (in kJ/mol) for the solution process. Assume that the specific heat of the solution is the same as that of pure water.

15. When a 4.25 g sample of solid ammonium nitrate dissolves in 60 g of water in a coffee-cup calorimeter, the temperature drops from 22.0°C to 16.9 °C. Calculate the change in enthalpy for the solution process and compare your calculated value to the one given in the chart on the previous page. Assume that the specific heat of the solution is that same as that of pure water. (answer = 25.8 kJ/mol)

16. Calculate the energy change (in J) if 3.65 g of magnesium sulfate heptahydrate is dissolved. (answer=239 J)

17. Calculate the energy (in J) absorbed when 10.3 g of barium chloride dihydrate is dissolved. (answer = 868 J)

18. A student is unsure if a bottle marked "Barium Chloride" is in anhydrous or hydrated form. What simple test could be performed to find out?

19. What mass of ammonium nitrate was dissolved if 8.93 kJ of energy is absorbed when the solution was made? (answer = 27.02 g ammonium nitrate)

20. Calculate ΔHsoln for copper (II) sulfate pentahydrate if 120.5 J of energy is absorbed when 2.50 g is dissolved to make a solution? (answer = 12,036 J/mol)

21. 5.23 g of sodium thiosulphate pentahydrate (Na2S2O3 · 5H2O) is dissolved in 500.0 mL of water. If the initial temperature of the water was 22.8(C, what will be the final temperature of the solution? State two assumptions you make in answering this question.

22. When sulfuric acid dissolves in water, a great deal of heat is given off. To measure it, 175 g of water was placed in a coffee-cup calorimeter and chilled to 10oC. Then 49.0 g of pure sulfuric acid, also at 10.0oC, was added and the mixture was quickly stirred. The temperature rose rapidly to 14.9oC.

a. Calculate q for the formation of this solution.

b. Calculate the molar enthalpy of solution ΔHsoln in kilojoules per mole of H2SO4.

23. Calculate the mass of anhydrous barium chloride dissolved in a 250.0 mL solution if, when added, the temperature of the mixture increased 3.79(C.

LAB 1: Enthalpy of Solution Lab written in lab notebook (omit sections II, III, IV, and VIII)

1. Measure approximately 1.0 grams of the salt (ionic compound) and record the actual mass.

2. Measure approximately 15.0 mL of distilled water with a graduated cylinder. Place the water into the cup. Add the lid and measure the temperature.

3. Add the salt to the water and observe the temperature until it hits the highest/lowest temperature.

4. Repeat for second salt.

5. Clean-up: pour solutions in waste beaker in fume hood and wash apparatus. Return materials.

***Section VII: calculate the enthalpy of solution for each salt in kJ/mol. Show all work.

LAB 2: Enthalpy of Phase Change Lab written on next page in lab notebook (omit sections II, III, IV, and VIII)

1. Fill two large beakers about ¾ full with tap water. Place one beaker with water on a hot plate and start heating the water. Leave the other beaker at room temperature.

2. Use a weigh boat to weigh about 5 grams of solid lauric acid (C12H24O2). Record the actual mass. Place the lauric acid into the test tube. Minimize the amount of lauric acid on the sides of the test tube.

3. Place test tube in the hot water bath (large beaker with hot water) and heat until the lauric acid is completely melted. While you’re waiting, continue with the next step.

4. Set up the temperature probe by selecting Sensor, then Data Collection, and changing the Rate to 20 and the Length to 15 minutes. Hit ‘ok’.

5. Once the lauric acid is completely melted, carefully remove the test tube from the hot water. Insert the temperature probe and press the arrow on the lab quest to start the data collection.

NOTE: make sure the thermometer is not touching the bottom of the test tube.

6. Once the 15 minutes have passed and the data collection has stopped, return the test tube to the hot water to melt it so that you can remove the temperature probe.

7. Clean up – pour the melted lauric acid into the waste container in the fume hood. Thoroughly clean out test tube. Return all supplies except hot plate. Turn off hot plate but do not unplug.

***Section VII:

a) Sketch the graph from the lab quest. Include a title and labels for the axes. Indicate where the solid is being cooled, where the liquid is being cooled, the phase change, and the freezing point.

b) How would this graph change if twice the mass of lauric acid were used?

Enthalpies of Phase Changes and Heats of Phase Changes

Vocab: fusion, vaporization, enthalpy of fusion, enthalpy of vaporization, heat of fusion, feat of vaporization

Objectives (students will be able to):

➢ define vocab terms

➢ use the LoCE to set up equations and calculate an unknown variable (SHC/temp/mass/enthalpy)

➢ understand and describe heating and cooling curbes

|Substance |Chemical Formula|Melting point (°C) |ΔHfus (kJ/mol) |Boiling point (°C) |ΔHvap (kJ/mol) |Specific Heat (J/g°C) |

|Water (liquid) |H2O |0.06 |6.01 |99.4 |40.7 |4.184 |

|Water (ice) |H2O |0.06 |6.01 | | |2.1 |

|Nitrogen |N2 |-209.7 |0.72 |-195.9 |5.58 |1.040 |

Notes/Problem Workspace:

How many kJ of heat are required to melt a 10.0 g popsicle at 0.06°C? (Assume the popsicle has identical properties to water.) Is this process endothermic or exothermic? (answer =3.34 kJ)

How much energy is required to convert 2500.0 grams of nitrogen gas at 22°C to liquid nitrogen at -195.9°C? Is this process endothermic or exothermic? (answer = -1065 kJ)

[pic]

24. How many grams of ice at 0.06°C could be melted by 0.400 kJ of heat? Is this process endothermic or exothermic? (answer = 1.20 g)

25. How much energy is required to cool 475.0 g of steam from 150°C to 101°C? Is this process endothermic or exothermic? (-39.6 kJ)

26. How much energy is required to convert 99.50 g of ethanol at 50.0°C to vapor at 78.5°C? Is this process endothermic or exothermic? (answer = 101 kJ)

27. An Olympic sized swimming pool holds 2,452,947 liters of water.

a. If this pool is filled with water from the public water supply in April when tap water is usually 16.00°C, how much heat energy is required to heat the pool to the comfortable temperature of 25.00°C? (answer = 9.233x1010 J)

b. Is this process endothermic or exothermic?

c. If the electric company charges $0.07710 for every 3.600x106 joules of electricity used (and we assume that all electricity was converted to heat energy since that is what we are using it for), how much will it cost to heat the pool to 25.00°C? (answer = $1,977)

28. How much energy is required to convert 825.0 grams of liquid nitrogen at -196.0°C to gas at 20.0°C? Is this process endothermic or exothermic? (answer = 350.0 kJ)

Hess’ Law

See for more practice

Vocab: Hess’ Law

Objectives (students will be able to):

➢ define vocab terms

➢ use Hess’ Law to calculate enthalpies

Notes/Problem Workspace:

MAIN RULE:

Strategies:

1.

2.

DO NOT STRAY FROM THESE STRATEGIES!!! Do only strategy 1 until you can’t anymore, and then switch to strategy 2.

Calculate the enthalpy of the reaction:

GOAL EQUATION: 3H2(g) + O3(g) ( 3H2O(g) ∆H = ?

Use the following enthalpies of reactions:

2H2(g) + O2(g) ( 2H2O(g) ∆H = - 483.6 kJ/mol rxn

3O2(g) ( 2O3(g) ∆H = 284.6 kJ/mol rxn

(answer = -867 kJ/mol rxn)

Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:

Goal equation: PCl5(g)  →  PCl3(g)  +  Cl2(g)

Use the following information:

P4(s)  +  6Cl2(g)  →  4PCl3(g)            ΔH = -2439 kJ/mol rxn

4PCl5(g)  →  P4(s)  +  10Cl 2(g)         ΔH = 3438 kJ/mol rxn

(answer = 249.8 kJ/mol rxn)

29. From the following enthalpies of reaction, calculate:

C2H4(g) + 6F2(g) ( 2CF4(g) + 4HF(g) ∆H = ?

H2 + F2(g) ( 2HF(g) ∆H = -537 kJ/mol rxn

C(s) + 2F2(g) ( CF4(g) ∆H = -680 kJ/mol rxn

2C(s) + 2H2(g) ( C2H4(g) ∆H = 52.3 kJ/mol rxn

answer = -2486 kJ/mol rxn

30. Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:

2CO2(g)  +  H2O(g)  →  C 2H2(g) +  5/2O2(g)

C2H2(g) + 2H2(g)  →  C2H6(g)                              ΔH  =-94.5 kJ/mol rxn

H2O(g)  →  H2(g) + 1/2O2 (g)                               ΔH  =71.2 kJ/mol rxn

C2H6(g) +  7/2O2(g)  →  2CO2(g)  +  3H2O(g)     ΔH  =-283 kJ/mol rxn

answer = 235 kJ/mol rxn

31. Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:

N2H4(l)  +  H2(g)  →  2NH3(g)

N2H4(l)  +  CH4O(l)  →  CH2O(g)  +  N2(g)  +  3H2 (g)         ΔH = -37 kJ/mol rxn

N2(g)  +  3H2(g)  →  2NH 3(g)                                                ΔH = -46 kJ/mol rxn

CH4O(l)  →  CH2O(g) +  H 2(g)                                              ΔH = -65 kJ/mol rxn

answer = -18 kJ/mol rxn

32. Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:

H2SO4(l)  →  SO3(g)  +  H2O(g)

H2S(g)  +  2O2(g)  →  H2SO4(l)                                  ΔH = -235.5 kJ/mol rxn

H2S(g)  +  2O2(g)  →  SO 3(g)  +  H2O(l)                    ΔH = -207 kJ/mol rxn

H2O(l)  →  H2O(g)                                                      ΔH = 44 kJ/mol rxn

answer = 72 kJ/mol rxn

33. Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:

C2H2(g) +  5/2O2(g)  →  2CO2(g)  +  H2O(g)

Use the following information:

C2H6(g)  →  C2H 2(g) + 2H2(g)                              ΔH = 283.5 kJ/mol rxn

H2(g) + 1/2O2(g)  →  H2O(g)                                ΔH = -213.7 kJ/mol rxn

2CO2(g)  +  3H2O(g)  →  C2H6(g) +  7/2O2(g)      ΔH = 849 kJ/mol rxn

answer = -705 kJ/mol rxn

34. Find the ΔH for the reaction below, given the following reactions and subsequent ΔH values:

Zn(s) +  1/8S8(s)  +  2O2(g)  →  ZnSO4(s)

Use the following information:

Zn(s) + 1/8S8(s)   →   ZnS(s)                                   ΔH = -183.92 kJ/mol rxn

2ZnS(s) + 3O2(g)   →   2ZnO(s) + 2SO2(g)            ΔH = -927.54 kJ/mol rxn

2SO2(g) + O2(g)   →   2SO3(g)                              ΔH = -196.04 kJ/mol rxn

ZnO(s) + SO3(g)   →   ZnSO4 (s)                           ΔH = -230.32 kJ/mol rxn

answer = -976.03 kJ/mol rxn

35. Calculate the ΔH for the reaction below:

N2H4(l) + O2(g) ( N2(g) + 2H2O(l)

Use the following information:

2NH3(g) + 3N2O(g) ( 4N2(g) + 3H2O(l) ΔH = -1010 kJ/mol rxn

N2O(g) + 3H2 ( N2H4(l) + H2O(l) ΔH = -317 kJ/mol rxn

2NH3(g) + ½ O2(g) ( N2H4(l) + H2O(l) ΔH = -143 kJ/mol rxn

H2(g) + ½ O2(g) ( H20(l) ΔH = -286 kJ/mol rxn

answer = -622.5 kJ/mol rxn

Enthalpies of Formation

Vocab: enthalpy of formation (ΔHf)

Objectives (students will be able to):

➢ define vocab terms

➢ Use the equation: ΔHºrxn = ΣnΔHºf(products) - ΣnΔHºf(reactants)

➢ use the LoCE to set up equations and calculate an unknown variable (SHC/temp/mass/enthalpy)

Notes/Problem Workspace:

The following is known as a thermite reaction: 2Al(s) + Fe2O3(s) ( Al2O3(s) + 2Fe(s). This highly exothermic reaction is used for welding massive units, such as propellers for large ships. Using enthalpies of formation, calculate ∆H° for this reaction. (answer = -850 kJ/mol rxn)

36. Complete combustion of 1 mol of acetone, C3H6O, results in the liberation of 1790 kJ.

a. Write the balanced chemical equation for the combustion of 1 mole of acetone.

b. Calculate the enthalpy of formation of acetone. Assume products are in the gas phase. (answer = -118 kJ/mol)

37. Calculate the standard enthalpy change for each of the following reactions:

a. 2SO2(g) + O2(g) ( 2SO3(g)

b. Mg(OH)2(s) ( MgO(s) + H2O(l)

c. 4FeO(s) + O2(g) ( 2Fe2O3(s)

d. SiCl4(l) + 2H2O(l) ( SiO2(s) + 4HCl(g)

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TIPS

1. When using the LoCE, begin by writing the -qlost=qgained (LoCE) equation (it’s like a map to the problem)

TIP

2. When problems look long and hard, start by writing what you’re looking for. Write the units your unknown should have. Then, identify which side of the LoCE equation it should be in.

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