Chapter 6: Thermochemistry 143
chapter 6
Thermochemistry
Chapter Terms and Definitions
Numbers in parentheses after definitions give the text sections in which the terms are explained. Starred terms are italicized in the text. Where a term does not fall directly under a text section heading, additional information is given for you to locate it.
thermodynamics* science of the relationships between heat and other forms of energy (6.1, introductory section)
thermochemistry* one area of thermodynamics; study of the quantity of heat absorbed or evolved by chemical reactions (6.1, introductory section)
energy potential or capacity to move matter (6.1)
kinetic energy (Ek) energy associated with an object by virtue of its motion (6.1)
joule (J) SI unit of energy, kg ∙ m2/s2 (6.1)
watt* measure of quantity of energy used per unit time; 1 J/s (6.1)
calorie (cal) non-SI unit of energy commonly used by chemists; originally defined as the amount of energy required to raise the temperature of one gram of water by one degree Celsius; 1 cal = 4.184 J (exact definition) (6.1)
potential energy (Ep) an object’s energy because of its position in a field of force (6.1)
internal energy (U) sum of kinetic and potential energies of the particles making up a substance (6.1)
Etot* total energy of a substance; sum of the kinetic, potential, and internal energies of the substance (6.1)
law of conservation of energy energy may be converted from one form to another, but the total quantity of energy remains constant (6.1)
thermodynamic system (or system) substance or mixture of substances under study in which a physical or chemical change occurs (6.2)
surroundings everything in the vicinity of a thermodynamic system (6.2)
heat (q) energy that flows between system and surroundings because of a difference in temperature between the thermodynamic system and its surroundings (6.2)
thermal equilibrium* state in which energy does not flow as heat between system and surroundings; temperature equality (6.2)
heat of reaction value of q required, at a given temperature, to return a system to the given temperature at the completion of the reaction (6.2)
exothermic process chemical reaction or physical change in which heat is evolved (q is negative) (6.2)
endothermic process chemical reaction or physical change in which heat is absorbed (q is positive) (6.2)
qp* heat of reaction at constant pressure (6.3)
enthalpy (H) extensive property of a substance used to obtain the heat absorbed or evolved in a chemical reaction (6.3)
state function property of a system that depends only on its present state, which is determined by variables such as temperature and pressure and is independent of any previous history of the system (6.3)
enthalpy of reaction (ΔH) change in enthalpy for a reaction at a given temperature and pressure; equals the heat of reaction at constant pressure (6.3)
enthalpy diagram* pictorial representation of the enthalpy change for a reaction (6.3)
pressure–volume work* energy required by a system to change volume against the constant pressure of the atmosphere (6.3)
thermochemical equation chemical equation for a reaction (including phase labels) in which the equation is given a molar interpretation, and the enthalpy of reaction for these molar amounts is written directly after the equation (6.4)
heat capacity (C) quantity of heat needed to raise the temperature of the sample of substance one degree Celsius (or one kelvin) (6.6)
specific heat capacity (specific heat) quantity of heat required to raise the temperature of one gram of a substance by one degree Celsius (or one kelvin) at constant pressure (6.6)
calorimeter device used to measure the heat absorbed or evolved during a physical or chemical change (6.6)
bomb calorimeter* calorimeter used for reactions involving gases (6.6)
Hess’s law of heat summation for a chemical equation that can be written as the sum of two or more steps, the enthalpy change for the overall equation equals the sum of the enthalpy changes for the individual steps (6.7)
standard state standard thermodynamic conditions chosen for substances when listing or comparing thermodynamic data: 1 atm pressure and the specified temperature (usually 25°C) (6.8)
standard enthalpy of reaction (ΔH°)* enthalpy change for a reaction in which reactants in their standard states yield products in their standard states (6.8)
allotrope one of two or more distinct forms of an element in the same physical state (6.8)
reference form stablest form (physical state and allotrope) of the element under standard thermodynamic conditions (6.8)
standard enthalpy of formation (standard heat of formation) () enthalpy change for the formation of one mole of a substance in its standard state from its elements in their reference forms and in their standard states (6.8)
fuel* any substance that is burned or similarly reacted to provide heat and other forms of energy (6.9)
Chapter Diagnostic Test
1. How much heat is produced when 8.95 g C2H5OH is burned in a constant-pressure system? The equation and enthalpy of reaction are
C2H5OH(l) + 3O2(g) ® 2CO2(g) + 3H2O(l) ΔH = -1.367 ´ 103 kJ
2. Calculate the enthalpy change at 298 K for the reaction
Ni(s) + 4CO(g) ® Ni(CO)4(g)
Heats of formation at 298 K are for CO = -110.5 kJ/mol
for Ni(CO)4 = -605 kJ/mol
3. Calculate ΔH° at 298 K for the reaction C (graphite) + CO2(g) ® 2CO(g)
using the following ΔH° data at 298 K:
H2(g) + CO(g) ® C (graphite) + H2O(g) / ΔH° = -131.38 kJFeO(s) + H2(g) ® Fe(s) + H2O(g) / ΔH° = 24.69 kJ
FeO(s) + CO(g) ® Fe(s) + CO2(g) / ΔH° = -16.32 kJ
4. What is the kinetic energy of an oxygen molecule traveling at a speed of 479 m/s in a tank at 21°C? (Hint: Use Avogadro’s number and the molar mass of O2 to get the actual mass of an oxygen molecule.)
5. For the reaction H2S(g) + 4H2O2(l) ® H2SO4(l) + 4H2O(l), ΔH° is -1.186 ´ 103 kJ. The enthalpy change per mole of H2O2 is
a. 1.301 ´ 102 kJ.
b. 4.742 ´ 103 kJ.
c. 2.965 ´ 102 kJ.
d. -4.741 ´ 103 kJ.
e. none of the above.
6. Indicate whether each of the following statements is true or false. If a statement is false, change it so that it is true.
a. The enthalpy of reaction is independent of the exact state of the reactants or products. True/False:______
______.
b. If the enthalpy of reaction for N2(g) + O2(g) ® 2NO(g) is 180.5 kJ, then the enthalpy of reaction for ½N2(g) + ½O2(g) ® NO(g) is 90.25 kJ. True/False:
______
______.
c. A calorimeter is a useful apparatus for determining heats of reaction. True/False:
______
______.
d. For reactions involving gases and carried out at constant pressure, ΔH = qp. True/False:
______
______.
e. Hess’s law permits the calculation of ΔH° values for reactions from values for reactants and products and the calculation of ΔH° values for hypothetical reactions. True/False: ______
______.
f. The value for an element is always zero. True/False: ______
______.
7. Use heat of formation data in Appendix C in the text to calculate the enthalpy of the transition from the liquid to the gaseous state for 1 mole of HCN. Report the answer in kilojoules and kilocalories.
8. Determine the enthalpy of ionization for Cs(g) using all the thermodynamic data below. The equation is Cs(g) Cs+(g) + e-(g).
F(g)+ e-(g) ® F-(g) / ΔH = -336 kJF2(g) ® 2F(g) / ΔH = 158 kJ
Cs(s) ® Cs(g) / ΔH = 78 kJ
Cs(s) + ½F2(g) ® CsF(s) / ΔH = -555 kJ
CsF(s) ® Cs+(g) + F-(g) / ΔH = 757 kJ
9. When 2.89 g N2H4(g) is combusted in a constant-pressure calorimeter containing exactly 1000 g of water, a temperature increase of 6.68°C is observed. The heat capacity of the calorimeter is 1.00 kJ/°C, and the specific heat of water is 4.184 J/(g ∙ °C). All products are gaseous. Determine the enthalpy change per mole of N2H4 combusted.
10. A 17.9-g sample of an unknown metal was heated to 48.31°C. It was then added to 28.05 g of water in an insulated cup. The water temperature rose from 21.04 to 23.98°C. What is the specific heat of the metal?
11. Ethane gas, C2H6, burns in oxygen to form carbon dioxide gas, CO2, and gaseous water. For each mole of ethane burned, 1.60 kJ of heat is evolved at constant pressure. Write the thermochemical equation for this reaction, including labels for the states of all reactants and products.
Answers to Chapter Diagnostic Test
If you missed an answer, study the text section and problem-solving skill (PS Sk.) given in parentheses after the answer.
1. -2.66 ´ 102 kJ (6.5, PS Sk. 4)
2. -163 kJ (6.8, PS Sk. 9)
3. 172.39 kJ (6.7, PS Sk. 7)
4. 6.10 ´ 10-21 J/molecule (6.1, PS Sk. 1)
5. e (6.4, PS Sk. 3)
6.
a. False. The enthalpy of reaction depends on the exact state of the reactants or products. (6.4)
b. True. (6.4, PS Sk. 3)
c. True. (6.6)
d. True. (6.3)
e. True. (6.7)
f. False. The value for an element is always zero when the state of the element is the form of the element that exists at standard conditions. (6.8)
7. 3.0 ´ 101 kJ; 7.2 kcal (6.8, PS Sk. 8)
8. 381 kJ (6.7, PS Sk. 7)
9. -383 kJ (6.6, PS Sk. 6)
10. 0.792 J/(g ∙ °C) (6.6, PS Sk. 5)
11. C2H6(g) + O2(g) ® 2CO2(g) + 3H2O(g), ΔH = -1.60 kJ (6.4, PS Sk. 2)
Summary of Chapter Topics
Students find thermochemistry one of the more difficult topics in chemistry. This is said not to scare you but to assure you that if you find yourself really scratching your head, you are not alone. It is also said to let you know that this material is going to take a great deal of time and study to master. Your text presents the subject quite well, but you probably will need to read it over several times before the concepts begin to make sense. One of the biggest stumbling blocks for students is the arithmetic signs (+ and -) that go with almost every term. If you memorize the sign conventions stressed in the text, you will find things considerably easier. If you are stuck in your thinking or in an exercise or problem, go back and review the sign conventions to see if your error is there.
6.1 Energy and Its Units
Learning Objectives
· Define energy, kinetic energy, potential energy, and internal energy.
· Define the SI unit of energy joule, as well as the common unit of energy calorie.
· Calculate the kinetic energy of a moving object. (Example 6.1)
· State the law of conservation of energy.
Problem-Solving Skill
1. Calculating kinetic energy. Given the mass and speed of an object, calculate the kinetic energy (Example 6.1).
Energy has many forms. Some of these are electromagnetic energy (such as light, heat, and x rays), electrical energy, sound energy, gravitational potential energy, elastic potential energy, and chemical potential energy. The forms of energy we will be concerned with in this course are electromagnetic energy, electrical energy, and chemical potential energy.
Exercise 6.1
An electron, whose mass is 9.11 ´ 10-31 kg, is accelerated by a positive charge to a speed of 5.0 ´ 106 m/s. What is the kinetic energy of the electron in joules? in calories?
Known: Ek = ½mv2; 1 J = 1 kg ∙ m2/s2; 1 cal = 4.184 J
Solution:
Ek = ´ 9.11 ´ 10–31 kg ´ (5.0 ´ 106 m/s)2 ´
= 1.1 ´ 10–17 J
Ek = 1.14 ´ 10–17 J ´ = 2.7 ´ 10–18 cal
6.2 Heat of Reaction
Heat is a difficult term to understand. The word heat makes a good verb but a poor noun. Heat is energy in transit from a hotter object to a colder one. Objects do not possess heat. They possess energy that can be transferred as heat. Once the energy arrives at its destination, it is absorbed and is no longer called heat.
In an exothermic reaction, the reactants are always at a higher state of enthalpy than are the products. Heat is given off, and the heat of reaction is negative (-). In an endothermic reaction, it is just the reverse. The reactants are at a lower state of enthalpy than are the products. Energy must be added to the reactants to get the products. Thus the heat of the reaction is positive (+). The following diagrams illustrate these concepts.
Exothermic Reaction / Endothermic ReactionLearning Objectives
· Define a thermodynamic system and its surroundings.
· Define heat and heat of reaction.
· Distinguish between an exothermic process and an endothermic process.
Exercise 6.2
Ammonia burns in the presence of a platinum catalyst to give nitric oxide, NO.
4NH3(g) + 5O2(g) 4NO(g) + 6H2O(l)
In an experiment, 4 mol NH3 is burned and evolves 1170 kJ of heat. Is the reaction endothermic or exothermic? What is the value of q?
Wanted: whether reaction is endothermic or exothermic; value of q
Given: 1170 kJ of heat evolves.
Known: definitions of two terms; sign is + if heat is absorbed.
Solution: Reaction is exothermic; q = -1170 kJ.
6.3 Enthalpy and Enthalpy Change
Learning Objectives
· Define enthalpy and enthalpy of reaction.
· Explain how the terms enthalpy of reaction and heat of reaction are related.
· Explain how enthalpy and internal energy are related.
6.4 Thermochemical Equations
Learning Objectives
· Define a thermochemical equation.
· Write a thermochemical equation given pertinent information. (Example 6.2)
· Learn the two rules for manipulating (reversing and multiplying) thermochemical equations.