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Thermochemistry

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Dr. Kittisak ChoojunDepartment of Chemistry,Faculty of Science, KMITL

Ref: Ralph H. Petrucci, F. ed. Geoffrey Herring, Jeffry D. Madura, Carey Bissonnette. General Chemistry: Principles and Modern Applications. Prentice Hall 2010, English 10th ed. Theodore E. Brown, H. Eugene H LeMay, Bruce E. Bursten, Catherine Murphy, Patrick Woodward. Chemistry: The Central Science. Prentice Hall 2011, English 12th ed. Raymond Chang and Jason Overby. General Chemistry: The Essential Concepts. McGraw-Hill Science/Engineering/Math 2010, 6th

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HeatHeat is energy transferred between a system and its surroundings as a result of a temperature difference. Energy that passes from a warmer body (with a higher temperature) to a colder body (with a lower temperature) is transferred as heat. During the process of melting, the temperature remains constant as a thermal energy transfer (heat) is used to overcome the forces holding the solid together. Aprocessoccurring at a constant temperature is said to be isothermal.

It is reasonable to expect that the quantity of heat, q, required to change the temperature of a substance depends on 1. how much the temperature is to be changed 2. the quantity of substance 3. the nature of the substance (type of atoms or molecules)

The quantity of heat required to change the temperature of one gram of water by one degree Celsius has been called the calorie (cal). The calorie is a small unit of energy, and the unit kilocalorie (kcal) has also been widely used.

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Heat

The quantity of heat required to change the temperature of a system by onedegree is called the heat capacity of the system. If the system is a mole of substance, the term molar heat capacity is applicable.If the system is one gram of substance, the applicable term is specific heat capacity, or more commonly, specific heat (sp ht).

Specific Heat of Water

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Heat

The law of conservation of energy: In interactions between a system and its surroundings, the total energy remains constant energy is neither created nor destroyed.

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HeatDetermining a Specific Heat from Experimental Data

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Specific Heats of Some Substances

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The specific heat of octane, C8H18(l), is 2.22 J/g.K. (a) How many J of heat are needed to raise the temperature of 80.0 g of octane

from 10.0 °C to 25.0 °C? (b) Which will require more heat, increasing the temperature of 1 mol of C8H18(l)

by a certain amount or increasing the temperature of 1 mol of H2O(l) by the same amount?

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Heats of Reaction and CalorimetryA heat of reaction, is the quantity of heat, qrxn, exchanged between a system and its surroundings when a chemical reaction occurs within the system at constant temperature.If a reaction occurs in an isolated system, that is, one that exchanges no matter or energy with its surroundings, the reaction produces a change in the thermal energy of the system the temperature either increases or decreases.

Heats of reaction are experimentally determined in a calorimeter, a device for measuring quantities of heat. We will consider two types of calorimeters in this section, and we will treat both of them as isolated systems.

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Heats of Reaction and CalorimetryBomb Calorimetry

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A 2.200-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 heat of combustion per gram of quinone? Per mole of quinone?

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WorkWork involved in the expansion or compression of gases is called pressure volume work. Pressure volume, or P V, work is the type of work performed by explosives and by the gases formed in the combustion of gasoline in an automobile engine.

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The First Law of Thermodynamics

Internal energy, U, is the total energy (both kinetic and potential) in a system, including translational kinetic energy of molecules, the energy associated with molecular rotations and vibrations, the energy stored in chemical bonds and intermolecular attractions, and the energy associated with electrons in atoms.

the law of conservation of energy

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Functions of StateAny property that has a unique value for a specified state of a system is said to be a function of state, or a state function. Thus, the value of a function of state depends on the state of the system, and not on how that state was established. The internal energy of a system is a function of state.

Unlike internal energy and changes in internal energy, heat (q) and work (w) are not functions of state.

Any state function depends only on the present state of the system and not on the path by which the system arrived at that state.

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Heat of Reaction: ∆U and ∆H

Bomb Calorimeter Constant Volume

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Heat of Reaction: ∆U and ∆HThe usual other way is in beakers, flasks, and other containers open to the atmosphere and under the constant pressure of the atmosphere.

The enthalpy, H, is the sum of the internal energy and the pressure volume product of a system.

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Indicate the sign of the enthalpy change, ∆H, in these processes carried out under atmospheric pressure and indicate whether each process is endothermic or exothermic: (a) An ice cube melts (b) 1 g of butane (C4H10) is combusted in sufficient oxygen to give complete

combustion to CO2 and H2O.

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Enthalpy Change (∆H) Accompanying a Change in State of Matter

Standard States and Standard Enthalpy Changes

The measured enthalpy change for a reaction has a unique value only if the initialstate (reactants) and final state (products) are precisely described. The standard state of a solid or liquid substance is the pure element or compound at a pressure of 1 bar * and at the temperature of interest. For a gas, the standard state is the pure gas behaving as an (hypothetical) ideal gasat a pressure of 1 bar and the temperature of interest. Although temperature is not part of the definition of a standard state, it still must be specified in tabulated values of because depends on temperature. The values given in this text are all for 298.15 K (25 °C) unless otherwise stated.

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The complete combustion of ethanol, C2H5OH(l), to form H2O(g) and CO2(g) at constant pressure releases 1235 kJ of heat per mole of C2H5OH. (a) Write a balanced thermochemical equation for this reaction. (b) Draw an enthalpy diagram for the reaction.

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Ozone,O3(g),is a form of elemental oxygen that is important in the absorption of ultraviolet radiation in the stratosphere. It decomposes to O2(g) at room temperature and pressure according to the following reaction:

2O3(g) à 3O2(g) ∆H = -284.6 kJ (a) What is the enthalpy change for this reaction per mole of O3(g)? (b) Which has the higher enthalpy under these conditions, 2O3(g) or 3O2(g)?

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Exercises 199

ENTHALPY (sections 5.3 and 5.4)5.33 (a) Why is the change in enthalpy usually easier to measure

than the change in internal energy? (b) H is a state function,but q is not a state function. Explain. (c) For a given process atconstant pressure, ¢H is positive. Is the process endothermicor exothermic?

5.34 (a) Under what condition will the enthalpy change of aprocess equal the amount of heat transferred into or out of thesystem? (b) During a constant-pressure process, the systemreleases heat to the surroundings. Does the enthalpy of the sys-tem increase or decrease during the process? (c) In a constant-pressure process, . What can you conclude about ¢E, q,and w?

5.35 You are given ¢H for a process that occurs at constant pres-sure. What additional information do you need to determine¢E for the process?

5.36 Suppose that the gas-phase reaction were carried out in a constant-volume container at

constant temperature. Would the measured heat change repre-sent ¢H or ¢E? If there is a difference, which quantity is largerfor this reaction? Explain.

5.37 A gas is confined to a cylinder under constant atmosphericpressure, as illustrated in Figure 5.4. When the gas undergoes aparticular chemical reaction, it absorbs 824 J of heat from itssurroundings and has 0.65 kJ of P-V work done on it by itssurroundings. What are the values of ¢H and ¢E for thisprocess?

5.38 A gas is confined to a cylinder under constant atmosphericpressure, as illustrated in Figure 5.4. When 0.49 kJ of heat isadded to the gas, it expands and does 214 J of work on thesurroundings. What are the values of ¢H and ¢E for thisprocess?

5.39 The complete combustion of ethanol, C2H5OH(l), to formH2O(g) and CO2(g) at constant pressure releases 1235 kJ ofheat per mole of C2H5OH. (a) Write a balanced thermochem-ical equation for this reaction. (b) Draw an enthalpy diagramfor the reaction.

5.40 The decomposition of slaked lime, Ca(OH)2(s), into lime,CaO(s), and H2O(g) at constant pressure requires the additionof 109 kJ of heat per mole of Ca(OH)2. (a) Write a balancedthermochemical equation for the reaction. (b) Draw an en-thalpy diagram for the reaction.

5.41 Ozone, O3(g), is a form of elemental oxygen that is importantin the absorption of ultraviolet radiation in the stratosphere.It decomposes to O2(g) at room temperature and pressure ac-cording to the following reaction:

(a) What is the enthalpy change for this reaction per mole ofO3(g)? (b) Which has the higher enthalpy under these condi-tions, 2 O3(g) or 3 O2(g)?

5.42 Without referring to tables, predict which of the following hasthe higher enthalpy in each case: (a) 1 mol CO2(s) or 1 molCO2(g) at the same temperature, (b) 2 mol of hydrogen atomsor 1 mol of H2, (c) 1 mol H2(g) and 0.5 mol O2(g) at 25 °C or1 mol H2O(g) at 25 °C, (d) 1 mol N2(g) at 100 °C or 1 molN2(g) at 300 °C.

2 O3(g) ¡ 3 O2(g) ¢H = -284.6 kJ

2 NO2(g)2 NO(g) + O2(g) ¡

¢H = 0

5.43 Consider the following reaction:

(a) Is this reaction exothermic or endothermic? (b) Calculatethe amount of heat transferred when 3.55 g of Mg(s) reacts atconstant pressure. (c) How many grams of MgO are producedduring an enthalpy change of (d) How many kilo-joules of heat are absorbed when 40.3 g of MgO(s) is decom-posed into Mg(s) and O2(g) at constant pressure?

5.44 Consider the following reaction:

(a) Is this reaction exothermic or endothermic? (b) Calculatethe amount of heat transferred when 24.0 g of CH3OH(g) isdecomposed by this reaction at constant pressure. (c) For agiven sample of CH3OH, the enthalpy change during the reac-tion is 82.1 kJ. How many grams of methane gas are produced?(d) How many kilojoules of heat are released when 38.5 g ofCH4(g) reacts completely with O2(g) to form CH3OH(g) atconstant pressure?

5.45 When solutions containing silver ions and chloride ions aremixed, silver chloride precipitates:

(a) Calculate ¢H for production of 0.450 mol of AgCl by thisreaction. (b) Calculate ¢H for the production of 9.00 g ofAgCl. (c) Calculate ¢H when dis-solves in water.

5.46 At one time, a common means of forming small quantities ofoxygen gas in the laboratory was to heat KClO3:

For this reaction, calculate ¢H for the formation of (a) 1.36mol of O2 and (b) 10.4 g of KCl. (c) The decomposition ofKClO3 proceeds spontaneously when it is heated. Do youthink that the reverse reaction, the formation of KClO3 fromKCl and O2, is likely to be feasible under ordinary conditions?Explain your answer.

5.47 Consider the combustion of liquid methanol, CH3OH(l):

(a) What is the enthalpy change for the reverse reaction?(b) Balance the forward reaction with whole-number coeffi-cients. What is ¢H for the reaction represented by this equa-tion? (c) Which is more likely to be thermodynamicallyfavored, the forward reaction or the reverse reaction? (d) If thereaction were written to produce H2O(g) instead of H2O(l),would you expect the magnitude of ¢H to increase, decrease,or stay the same? Explain.

5.48 Consider the decomposition of liquid benzene, C6H6(l), togaseous acetylene, C2H2(g):

(a) What is the enthalpy change for the reverse reaction? (b)What is ¢H for the formation of 1 mol of acetylene? (c) Whichis more likely to be thermodynamically favored, the forwardreaction or the reverse reaction? (d) If C6H6(g) were con-sumed instead of C6H6(l), would you expect the magnitude of¢H to increase, decrease, or stay the same? Explain.

C6H6(l) ¡ 3 C2H2(g) ¢H = +630 kJ

¢H = -726.5 kJ

CH3OH(l) + 32 O2(g) ¡ CO2(g) + 2 H2O(l)

2 KClO3(s) ¡ 2 KCl(s) + 3 O2(g) ¢H = -89.4 kJ

9.25 * 10-4 mol of AgCl

Ag+(aq) + Cl-(aq) ¡ AgCl(s) ¢H = -65.5 kJ

2 CH3OH(g) ¡ 2 CH4(g) + O2(g) ¢H = +252.8 kJ

-234 kJ?

2 Mg(s) + O2(g) ¡ 2 MgO(s) ¢H = -1204 kJ

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The standard enthalpy of formation of a substance is the enthalpy change that occurs in the formation of one mole of the substance in the standard state from the reference forms of the elements in their standard states.

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ENTHALPIES OF REACTION

Because ∆H = Hfinal - Hinitial, the enthalpy change for a chemical reaction is given by

∆H = Hproducts - Hreactants

The enthalpy change that accompanies a reaction is called either the enthalpy of reaction or the heat of reaction and is sometimes written ∆Hrxn, where “rxn” is a commonly used abbreviation for “reaction.”

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Methanol (CH3OH) is used as a fuel in race cars. (a) Write a balanced equation for the combustion of liquid methanol in air. (b) Calculate the standard enthalpy change for the reac- tion, assuming H2O(g) as

a product. (c) Calculate the heat produced by combustion per liter of methanol. Methanol has

a density of 0.791 g/mL.(d) Calculate the mass of CO2 produced per kJ of heat emitted.

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A sample of a hydrocarbon is combusted completely in O2(g) to produce 21.83 g CO2(g), 4.47 g H2O(g), and 311 kJ of heat. (a) What is the mass of the hydrocarbon sample that was combusted? (b) What is the empirical formula of the hydrocarbon? (c) Calculate the value of ∆Hf ° per empirical-formula unit of the hydrocarbon.