CH 8 Thermochemistry Worksheet eboard 0708

Objectives:
1. Distinguish between system and surroundings.
2. Distinguish between heat energy and temperature. Name the SI units of temperature and thermal energy.
3. Calculate the heat content (thermal energy) of a substance using its specific heat, temperature and mass.
4. Solve calorimetry problems.
5. Name the energy associated with a change of state and calculate the energy required to change a given mass of substance from solid to liquid or liquid to gas.
6. Describe the changes in the microscopic structure of a substance during heating or cooling.
7. Explain the properties of gases, liquids, and solids in terms of the kinetic theory of matter.
8. Distinguish between endothermic and exothermic reactions.
9. Use thermochemical equations to describe energy changes in a chemical reaction.

Readings and Assignments:

Reading:
1. Principles of Heat Flow Section 8.1, pages 216-219
2. Measurement of Heat Flow Section 8.2, pages 219-222
3. Enthalpy Section 8.3, pages 222-223
4. Thermochemical Equations Sections 8.4, pages 223-226

Written Assignment
1. Principles of Heat Flow Packet, page 3, questions 1-10
2. Phase Changes Packet, page 4, questions 1-14
3. Measurement of Heat Flow Page 241, exercises # 2, 4, 6
*4. Enthalpy and Heat of Formation Page 242, exercise # 16, 20, 34; packet, page 7, questions 1-6 ; page 8, questions 1-5

Reference Tables

Specific Heat, Heat of Fusion, and Heat of Vaporization of Selected Compounds

Substance
Specific Heat, J/g  0C
Heat of Fusion, J/g
Heat of Vaporization, J/g

Aluminum, Al
0.9025
395.01
10742.6
Benzene, C6H6
1.74
126.15
394.9
Brass
0.376

Carbon, C
0.710

Copper, Cu
0.385
204.82
5308.6
Ethyl alcohol, CH3CH2OH
2.4194
108.68
877.8
Glass
0.664

Gold
0.129
62.98
1647.5
Ice
2.090
333.56
2257.2
Iron
0.4598
266.27
7440.4
Lead
0.1296
24.66
944.68
Magnesium
1.0157
371.60
6061.0
Mercury
0.1379
11.286
284.24
Silver
0.235

Steam
2.015
333.56
2257.2
Tin
0.2274

Titanium, Ti
0.5226

Water
4.180
333.56
2257.2
Zinc
0.388

Calorimetry(Heat)

1. The specific heat of aluminum is 0.895 J/g oC. Calculate the heat necessary to raise the temperature of 40.0 g of aluminum from -20.0 oC to 32.3 oC. Specific heat of aluminum is 0.215-cal/g oC.

2. Heated bricks or blocks of iron were used long ago to warm beds. A 1.49-kg block of iron heated to 155 oC would release how many joules of heat as it is cooled to 22 oC? The specific heat of iron is 0.4494 J/ g oC.

3. Dysprosium, Dy, was discovered in 1886. Its freezing point is 1400 oC and its boiling point is 2600 oC. If 2.0 x 103 joules of heat were added to heat 10.00 grams of dysprosium from its freezing point to its boiling point what is its specific heat?

4. Copper has a specific heat of 0.384 J/g oC. A 105-g sample is exposed to 15.2 kJ in an insulated container. How many degrees will the temperature of the copper sample increase?

Calorimetry (Heat Exchange)

5. Suppose a piece of gold with a mass of 21.5 g at temperature of 95.00 oC is dropped into an insulated calorimeter containing 125.0 g of water at 22.00 oC. What will be the final temperature of the water?

6. If a piece of iron with a mass of 7.86 g and a temperature of -25.0 oC is dropped in 35.0 cm3 of water at 21.0 oC, what will be the final temperature of the system?

7. Determine the final temperature of a system after 100.0 g of zinc at 95.0 oC is immersed into 50.0 g of water at 15.0 oC in a coffee cup calorimeter. Assume no loss of heat to the surroundings and to the calorimeter.

8. If a piece of silver (Cp = 0.2165 J/g x 0C) with mass of 14.16 g and temperature of 133.5 0C is dropped into 250.0 g of water at 17.20 0C, what will be the final temperature of the system?

9. A blacksmith heated an iron bar to 1445 oC. The blacksmith then tempered the metal by dropping it into 42 800 cm3 of water that had a temperature of 22 oC. The final temperature of the system was 45 oC. What was the mass of the bar?

10. A 752-cm3 sample of water was placed in a 1.00-kg aluminum pan. The initial temperature of the pan was 26 oC, and the final temperature of the system was 39 oC. What was the initial temperature of the water?

1. 1.87 x 103 Joules 2. 8.90 x 104 J 3. 0.1733 J/g oC 4. 377 oC
5. 22.4 oC 6. 19.9 oC 7. 27.5 oC
8. 17.54 0C 9. 6.5 x 103 g 10. Ti = 43 oC
Calorimetry and Phase Changes

1. How many calories are required to completely convert 5.0 g of ice at 0.0 oC into water vapor at 100 oC?

2. Calculate the heat evolved when 20.3 g of molten aluminum at its melting point of 658 oC is solidified and cooled to 25.0 oC. The enthalpy of fusion of aluminum is 94.5 cal/g.

3. How much heat is required to convert 100.0 g of ice at 0.00 oC to water vapor (steam) at 150.0 oC?

4. What will be the final temperature of the system when 3500 J of energy are added to 9.0 g of ice at 0.0 oC? Heat of fusion of ice is equal to 334 J/g.

5. How much heat is required to raise the temperature of 10.0 g of H2O(s) from -10.0 oC to H2O (g) at 150 oC?

6. If 100.0 g of pure water at 27 oC is placed into an insulated flask, how many grams of ice at 0.0 oC must be added to lower the temperature of the water to 5.0 oC? (Recall: energy lost = energy gained)

7. 50.0 g of Freon-11 (CClF3) changes from 4.0 oC to 29.0 oC. How much energy does it absorb? Freon -11 boils at 24.0 oC. Specific heat of liquid = 0.20 cal/ g oC; specific heat of vapor = 0.14 cal /g oC heat of vaporization = 43.0 cal/g.

Calorimetry and Phase Changes: Miscellaneous Problems

8. Haw many calories are required to heat each of the following from 15 oC to 65.0 oC? a. 1.0 g of water? b. 20.0- g of platinum? Specific heats: Pyres glass = 0.20 cal /g/; Pt = 0.032 cal/gK

9. A 45.0 g sample of an alloy was heated to 90.0 oC and then dropped into a beaker containing 82.0 g of water at 23.50 oC. The temperature of the water rose to a final 26.25 oC. What is the specific heat capacity of the alloy?

10. The melting point of a certain substance is 70.0 oC, its normal boiling point is 450.0 oC, its enthalpy of fusion is 30.0 cal/g, its enthalpy of vaporization is 45.0 cal/g, and its specific heat is 0.215 cal/g K. Calculate the heat required to convert 100.0 g of the substance from the solid state at 70.0 oC to vapor at 450.0 oC.

11. The specific heat of silver is 0.0565 cal/ g K. Assuming no loss of heat to the surrounding or to the container, calculate the final temperature when 100. G of silver at 40.0 oC is immersed in 60.0 g of water at 10.0 oC.

12. Determine the resulting temperature when 1.00 kg of ice at 0.0 oC is mixed with 9.00 kg of water at 50.0 oC. Heat of fusion of ice is 80.0 cal/g.

13. When 1.0 g of ammonium nitrate, NH4NO3, dissolves in 50.0 g of water, the temperature of the water drops from 25.0 oC to 23.32 oC. What is the molar heat of solution?

14. The combustion of natural gas (methane) yields 890.4 kJ/mole. How many grams of methane, CH4, would be needed to raise the temperature of 25.0 kg of water from 20.0 oC to 75.0 oC? Specific heat of water = 4.18 J/g oC.

1. 1.5 x 104 J 2. -4.68 kCal 3. 74.4 kCal 4. 13 oC
5. 7.49 x 103 cal 6. 27.5 g ice
8. 50. cal; 32.0 cal 9. 7.87 x 10-2 cal/g oC
10. 15.7 kCal 11. 12.1 oC 12. 37 oC
13. 6.72 cal/mol; 2.81 x 104 J/mol 14. 103. g of CH4

Phase Changes

Thermochemical Equations

Translate the following statements into thermochemical equations.

1. The heat of formation of Ag2O is -7.3 kcal/mol.

2. The heat of combustion of carbon is -94.1 kcal/mol.

3. The condensation of water vapor (H = 9.72 kcal).

4. The vaporization of benzene (H = 7.36 kcal/mol)

5. Using the thermochemical equation: CuO(s) + H2(g) ----> Cu(s) + H2O(l) H = -310.6 kcal, calculate H when a. 2.00 moles of CuO react c. 1.00 g of Cu is formed
b. 1.00 g of CuO reacts d. 40.0 L of H2 react at STP.

6. Using the thermochemical equation: CaO(s) + SO3(g) ----> CaSO4(s) H = -96.0 kcal, calculate the mass in grams of CaSO4 which must be formed to evolve 1.00 kcal of heat.

5. a. -52.4 kcal; b. -0.390 kcal; c. -0.490 kcal; d. -55.7 kcal 6. 1.42 g

Endothermic Reactions

Determine if the following reactions are endothermic or exothermic. 1. melting ice _____________________
2. boiling water _____________________
3. evaporating snow _____________________
4. burning coal _____________________
5. freezing water _____________________
6. condensing steam _____________________
7. forming ice crystals _____________________
8. exploding a firecracker _____________________

Enthalpy of Reactions

1. Which of the following are exothermic reactions? a. N2(g) + 3H2(g) → 2NH3(g) H = -22 kcal b. 2HgO (s) + 43.3 kcal → 2Hg(l) + O2(g) c. Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) + 27.6 kJ d. N2(g) + O2(g) → 2NO(g) H = +180.6 kJ

2. The reaction 2CO(g) + O2(g) → 2CO2(g) is exothermic. a. What is the sign of H for this reaction? b. What is the sign of H for the reverse reaction? c. Is the heat content of the product (2mol CO2) greater or less than that of the reactants ( 2 mol CO + 1 mol O2)?

3. Draw a heat content diagram for the combustion of acetylene, C2H2 (H = -310.6 kcal).

4. For the reaction N2(g) + O2(g) →→ 2NO(g) at 298 oC, H = +180.6 kJ. a. Which is greater, the total value of Hfo for the reactants or Hfo of the products? b. Calculate the Hfo of NO(g) . c. How does the strength of the bonds in the NO molecule compare with those in N2(g) and O2(g) molecules?

5. Given the equation

H2(g) +1/2O2(g) → H2O(g) Hr = -241.6 kJ a. How much heat is evolved in the formation of 2.50 moles of H2O(g)? b. How much heat is liberated when 1.75 moles of H2(g) burn in excess of oxygen? c. Write a thermochemical equation corresponding to the above reaction for the combustion of 2.00 moles of hydrogen. d. How much heat will be absorbed when 4.00 moles of H2O(g) will decompose in the reverse reaction? e. How much heat is evolved for each gram of hydrogen burned?

. ΔHfo = -20 kcal/mol C2H6 5. a. -606 kJ; b. -423 kJ; d. 966 kJ; e. -121 kJ Enthalpy - Review Activity

1. The specific heat of ethyl alcohol is 2.43 J/g oC. If the ethyl alcohol absorbs 17.85 J of heat when its temperature increases from 23.13 oC to 27.48 oC, what is the mass of the ethyl alcohol present? 1.69 g

2. Calculate the specific heat of n alloy if 17.98 J of heat must be absorbed to raise the temperature of a 6.00-g sample from 20.00 oC to 28.35 oC.

3. When 1.507 g of NH4Cl are dissolved in 100.00 g of water at 25.00 oC, the temperature drops to 23.98 oC. Calculate the heat absorbed by the water and the heat released during this reaction. -427 J; 427 J

4. A 1.000 g sample of ethanol, C2H5OH was burned in a coffee cup calorimeter containing 1220 g of water. The temperature of the water in the calorimeter increased by 5.252 oC . How much heat did the burning of the ethanol generate? What would be the enthalpy of combustion for 1 mole of ethanol? -26.81kJ; 1235 kJ

5. When a piece of sodium metal is added to a beaker, the temperature in the beaker goes from 25 oC to 42 oC. Is the reaction exothermic or endothermic? What is the sign of ΔH?

6. Consider the reaction between sodium metal and liquid water. It produces solid sodium hydroxide, hydrogen gas, and 281.8 kJ of heat energy. Write the thermochemical equation for this reaction.

7. When four moles of iron metal are burned in oxygen, iron(III) oxide is formed and 1644.4 kJ of energy are evolved. a. Write a thermochemical equation for this reaction. b. How much heat is evolved if 3.031g of iron are burned? b) 22.31 kJ

8. Calculate the heat of formation of sodium oxide using the Thermodynamic Tables (Table 8.3 and 8.4 in your textbook) and using the following thermochemical equation: Na2O(s) + H2O(l) → 2Na+1(aq) + 2OH-1(aq) ΔH = -237.5 kJ

e. need more information

Notes:

Thermochemistry: study of heat flow that accompanies chemical reactions.
System: part of the universe upon which attention of the scientist is focused.
Surroundings: all the parts of the universe that are outside the system.
Exothermic process: heat is released. q<0; heat is flowing out of the system.
Endothermic process: heat is absorbed; q>0; heat is flowing into the system.
Magnitude of heat flow depends on:
1. mass of the object
2. change of temperature of the object
3. specific heat of the substance the object is made off.

Q = mct

m, mass of the substance cp, specific heat of the substance
t, change in temperature

Heat: energy that flows from a system at higher temperature to a system at lower temperature: the sum total of the KE of the individual particles in a system; represented by Q.
Temperature: degree of hotness or coldness; measures the average KE of a sample (system).

Work: can be done on the system by the surrounding or by the system on the surrounding.
Energy: the ability to do work (the system has the ability to do work because it has stored energy).

State of a system: is described by giving its composition, temperature, and pressure.
(Ex. H2O(l) at 50 ºC and 1 atm)
State Properties: depend only on the state of the system, not upon the way the system reached that state.

ΔX = Xfinal – Xinitial

Units of Heat: 1 calorie – heat needed to raise the temperature of 1.0 g of water 1 ºC at 4 ºC . 1 Joule = force of 1 Newton*1 meter (kg m2/sec2) 1 calorie = 4.18 joules
Thermometer scales: Celsius, Kelvin, Fahrenheit

Measuring energy changes:
Calorimeter: devise used to measure the energy changes. It is used routinely to measure changes of energy during chemical reactions. Types: coffee-cup calorimeter; bomb-calorimeter.
Quantity of heat depends on:
The mass of the substance in the calorimeter
Mass of the calorimeter
Specific heat of the substance
Change in temperature.

Heat capacity: quantity of thermal energy needed to raise the temperature of a given mass of substance by 1 ºC. It is an extensive property. Units: Joules/ºC
Specific heat: heat needed to raise the temperature of 1.0 g of substance 1 ºC. Units: Joules/(g ºC). Intensive property.
Law of conservation of energy: Heat gained = Heat lost; Q = mcΔT Endo q<0; Exo q>0

Change of State:

Solid: highly ordered arrangement; strong intermolecular forces (ionic, molecular, metallic, network covalent).
Melting/Boiling Process: Solids and liquids expand when heated. Increased KE causes particles to vibrate. At certain point particles move far apart, interaction between molecules weakens, the highly structured arrangement of liquids are somewhat more loosely arrangement of liquids breaks downs → the system is either melting or boiling.
Freezing/Condensation: the KE of particles decreases, the particles move slower, because of the decreased distances the intermolecular forces become stronger and condensation/freezing occurs.

Melting Point: the temperature at which a substances changes from solid to liquid.
Freezing Point: the temperature at which a substance changes from liquid to solid.
Boiling Point: the temperature at which a substance changes from liquid to gas.
Condensation Point: the temperature at which substance changes from gas to liquid.

All pure liquids have definite freezing point.
All pure solids have a definite melting point.
Characteristic Properties: melting point, freezing point, boiling point, condensation point.
All of the above are intensive properties.

Heat of Fusion: heat required to melt 1.0 g of solid of a specific substance at its melting point. Intensive property.

Enthalpy: type of chemical energy referred to as “heat content” . It is a property determined by the state of the substance. It is a chemical energy.
Origin o enthalpy: position and motion of atoms, molecules, and subatomic particles.