Chemistry
Enthalpy of Reaction
The enthalpy of reaction is the heat change that occurs during a chemical reaction at constant pressure. It is represented by the symbol ΔH and can be exothermic (releasing heat) or endothermic (absorbing heat). The enthalpy of reaction is a key concept in understanding the energy changes associated with chemical reactions.
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12 Key excerpts on "Enthalpy of Reaction"
eBook - PDF- R.E. Hayes, J.P. Mmbaga(Authors)
- 2012(Publication Date)
- CRC Press(Publisher)
19 2 Thermodynamics of Chemical Reactions Thermodynamics was identified in Chapter 1 as one of the key pillars in the field of chemical reactor analysis. Indeed, few reactor designs can be successfully executed in the absence of comprehensive thermodynamic calculations or, at the least, making some key thermodynamic assumptions. Several areas of chemical engineering thermodynamics are of major importance in chemical reaction engineering. They include, but are not limited to, the following: 1. Enthalpy change of reaction and enthalpy changes owing to temperature changes . Knowledge of these factors enables the prediction of temperature changes in the reactor and the calculation of heating/cooling requirements. 2. Chemical reaction equilibrium calculations enable the prediction of the maximum theoretical conversion at a given temperature. 3. Vapor liquid equilibrium is important in systems with both gaseous and liquid spe-cies. In such systems, a combined chemical reaction and phase equilibrium may be present. 2.1 Basic Definitions The following sections give a brief introduction to the basic thermodynamic terms. A more detailed presentation of the enthalpy changes during reactions is then given. Finally, the calculation of reaction equilibrium is reviewed. 2.1.1 Open and Closed Systems Thermodynamics divides the universe into a system and its surroundings. The system is the portion of the universe (e.g., a chemical reactor) that is to be analyzed while the remain-der of the universe is the surroundings. The system boundary separates the system from the surroundings. The definition of the system is arbitrary, and depends simply on where one draws the boundary. Chemical reactors may be either open or closed systems. In a closed system, the mass of the total material in the reactor remains constant and no mass enters or leaves the system. The volume of a closed system may change, and energy in the form of work or heat can enter or leave the system.
eBook - PDF- Charles R. Russell(Author)
- 2013(Publication Date)
- Pergamon(Publisher)
Thermochemical equations include the chemical reaction and resulting changes in enthalpy or free energy for the reactants and products in their indicated states, as # 2 (g) + K> 2 (g) = H 2 0 (1) AH° = -68.317 kg-cal/g-mole Unless otherwise indicated, the reactants and products are at a pressure of one atmosphere and some standard temperature, T°, such as 298 °K. In the thermochemical equation above, energy has been released (exothermic) since the enthalpy of the product is less than that of the reactants. When the heat of reaction is measured in an open calorimeter at constant pressure and otherwise at standard conditions, the standard enthalpy change is AH p = ΔΗ° When the heat of reaction is measured in a closed vessel such as a combustion bomb calorimeter at constant volume, the change in internal energy, AU° 9 is determined since AH V = AU° These values are related by AH° = AU°+pAV Where only condensed phases (liquids and solids) are involved, the difference between the values of the enthalpy and internal energy changes is very small and may be neglected. For gaseous reactants and products, the values of pAV, assuming an ideal gas, is pA V = AnRT° The value of An is the net increase in the number of moles of gaseous products. Nearly two centuries ago studies of heats of reactions led to the conclusion that the energy required to decompose a compound into its elements exactly equals the negative of the energy required in forming the compound from its elements. Later it was estab- 142 ELEMENTS OF ENERGY CONVERSION lished that the amount of energy associated with the formation of a compound from its elements is the same whether the reaction takes place in a single step or in a series of steps. By writing the chemical equations with notations for the state of each material (s —solid; 1 —liquid; g —gas) and giving the corresponding stand-ard heats of reactions, a systematic procedure of adding and subtracting reactions has been established.
eBook - PDF- Young, William Vining, Roberta Day, Beatrice Botch(Authors)
- 2017(Publication Date)
- Cengage Learning EMEA(Publisher)
Vasilyev/Shutterstock.com Thermochemistry Unit Outline 10.1 Energy 10.2 Enthalpy 10.3 Energy, Temperature Changes, and Changes of State 10.4 Enthalpy Changes and Chemical Reactions 10.5 Hess’s Law 10.6 Standard Heats of Reaction In This Unit… This unit begins an exploration of thermochemistry, the study of the role that energy in the form of heat plays in chemical processes. We inves-tigate the energy changes that take place during phase changes and the chemical reactions you have studied previously and learn why some chemical reactions occur while others do not. In Electromagnetic Radiation and the Electronic Structure of the Atom (Unit 3), you stud-ied energy changes at the molecular level and the consequences those energy changes have on the properties of atoms and elements. 10 Copyright 2018 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. WCN 02-300 Unit 10 Thermochemistry 272 10.1 Energy 10.1a Energy and Energy Units Chemical reactions involve reactants undergoing chemical change to form new substances, products. reactants S products What is not apparent in the preceding equation is the role of energy in a reaction. For many reactions, energy, often in the form of heat, is absorbed—that is, it acts somewhat like a reactant. You might write an equation for those reactions that looks like this: energy 1 reactants S products In other reactions, energy is produced—that is, it acts like a product: reactants S products 1 energy In Chemistry, Matter on the Atomic Scale (Unit 1), we defined energy (the ability to do work) and work (the force involved in moving an object some distance). From a chemist’s point of view, energy is best viewed as the ability to cause change, and thermochemistry is the study of how energy in the form of heat is involved in chemical change.
eBook - PDFChemistry
Principles and Reactions
- William Masterton, Cecile Hurley(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s). Editorial review has deemed that any suppressed content does not materially affect the overall learning experience. Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 196 CHAPTER 8 Thermochemistry ▼ Figure 8.6a shows the enthalpy relationship between reactants and products for an exothermic reaction such as CH 4 ( g ) 1 2O 2 ( g ) 9: CO 2 ( g ) 1 2H 2 O( l ) D H , 0 Here, the products, 1 mol of CO 2 ( g ) and 2 mol of H 2 O( l ), have a lower enthalpy than the reactants, 1 mol of CH 4 ( g ) and 2 mol of O 2 ( g ). The decrease in enthalpy is the source of the heat evolved to the surroundings. ▲ Figure 8.6b shows the situation for an endothermic process such as H 2 O( s ) 9: H 2 O( l ) D H . 0 Liquid water has a higher enthalpy than ice, so heat must be transferred from the surroundings to melt the ice. In general, the following relations apply for reactions taking place at constant pressure. exothermic reaction: q 5 D H , 0 H products , H reactants endothermic reaction: q 5 D H . 0 H products . H reactants The enthalpy of a substance, like its volume, is a state property. A sample of one gram of liquid water at 25.00 8 C and 1 atm has a fixed enthalpy, H . In prac-tice, no attempt is made to determine absolute values of enthalpy. Instead, scien-tists deal with changes in enthalpy, which are readily determined. For the process 1.00 g H 2 O ( l, 25.00 8 C, 1 atm) 9: 1.00 g H 2 O ( l, 26.00 8 C, 1 atm) D H is 4.18 J because the specific heat of water is 4.18 J/g 8 C. 8-4 Thermochemical Equations A chemical equation that shows the enthalpy relation between products and reactants is called a thermochemical equation . This type of equation contains, at the right of the balanced chemical equation, the appropriate value and sign for D H .
eBook - PDFChemistry
The Molecular Nature of Matter
- James E. Brady, Neil D. Jespersen, Alison Hyslop(Authors)
- 2014(Publication Date)
- Wiley(Publisher)
Describe the assumptions and utility of thermochemical equations A balanced chemical equation that includes both the enthalpy change and the physical states of the substances is called a ther- mochemical equation. Coefficients in a thermochemical equa- tion represent mole quantities of reactants and products, so it is reasonable to have fractional coefficients such as 1 2 . Such equa- tions can be added, reversed (reversing also the sign of ∆H ), or multiplied by a constant multiplier (doing the same to ∆H ). If formulas are canceled or added, they must be of substances in identical physical states. The reference conditions for thermochemistry, called stan- dard conditions, are 25 °C and 1 bar of pressure. An enthalpy change measured under these conditions is called the standard Enthalpy of Reaction or the standard heat of reaction, given the symbol ∆H °. Use Hess's law to determine the enthalpy of a reaction Hess’s law of heat summation is possible because enthalpy is a state function. Values of ∆H ° can be determined by the manipu- lation of any combination of thermochemical equations that add up to the final net equation. The units for ∆H ° are generally joules or kilojoules. An enthalpy diagram provides a graphical description of the enthalpy changes for alternative paths from reactants to products. Determine and use standard heats of formation to solve problems When the enthalpy change is for the complete combustion of one mole of a pure substance under standard conditions in pure oxygen, ∆H ° is called the standard heat of combustion of the compound and is symbolized as ∆H c °. When the enthalpy change is for the formation of one mole of a substance under standard conditions from its elements in their standard states, ∆H ° is called the standard heat of formation of the compound and is symbolized as ∆H f °, (usually in units of kilojoules per mole, kJ mol -1 ).
eBook - PDF- Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
The opposite is true if the system is compressed. We now rewrite: ΔU = q + w under conditions of constant pressure as shown below, where q p is the heat of reaction at constant pressure. ΔU = q p - pΔV However, this equation is inconvenient as we need to know the value of ΔV if we wish to calculate ΔU. We now define a new thermodynamic function called enthalpy, which has the symbol H, as: H = U + pV Thus, under conditions of constant pressure: ΔH = ΔU + pΔV Substituting ΔU = q p - pΔV into this gives the following equation. ΔH = q p - pΔV + pΔV Therefore, ΔH = q p Thus, the heat of reaction at constant pressure is equal to Δ r H, in the same way that the heat of reaction at constant volume is equal to Δ r U. Like Δ r U, Δ r H is a state function, meaning that the enthalpy change for a reaction depends only on the initial and final states of the system. If the final enthalpy of the system is greater than the initial, the system has absorbed heat from its surroundings, so Δ r H will be positive and the reaction is said to be endothermic. Conversely, if the system has lost heat to the surroundings, the enthalpy of the system has decreased, so Δ r H will be negative and the reaction is said to be exothermic. The difference between the enthalpy change and the internal energy change for a reaction is pΔV. The difference between Δ r U and Δ r H can be very large for reactions that produce or consume gases, because these reactions can have very large volume changes. If a reaction involves only solids and liquids, though, the values of ΔV are tiny, so Δ r U and Δ r H for these reactions are nearly identical. FIGURE 8.12 A coffee cup calorimeter used to measure heats of reaction at constant pressure insulated cover reactants in solution two nested polystyrene cups stirrer thermometer We can obtain values of Δ r H using a very simple constant pressure calorimeter, dubbed the coffee cup calorimeter.
- Elias I. Franses(Author)
- 2014(Publication Date)
- Cambridge University Press(Publisher)
The initial state and fi nal state are taken to be at the same temperature T and pressure p . The only changes between the fi nal and initial states are due to the reaction (16.19) . The energy balance equation, Eq. (4.9) , with no changes in the kinetic or the potential energy is Δ U ¼ Q þ W : ð 16 : 24 Þ For reversible changes at a fi xed pressure p , W = − p Δ V . Then Eq. (16.24) becomes Δ U − ð − p Δ V Þ ¼ Q , or, since Δ p = 0, Δ H ¼ Q : ð 16 : 25 Þ input flow stream 1 mole of A 1 mole of B per unit time outflow stream 1 mole of C 1 mole of D per unit time Q (net heat input) Figure 16.2 A schematic representation of a continuous reactor operation at steady state. 16.5 Enthalpy changes of a chemical reaction 365 Thus, the “ heat of reaction ” equals Δ H also for a batch reactor per mole of A, which is used as the “ basis. ” We will now describe how to determine the so-called “ standard heats, ” or “ enthalpies ” (meaning enthalpy changes ), of a reaction, Δ H ( T ), at p = 1 atm. First, we will consider the case of T = T 0 = 298 K. For the example of Reaction (16.19) Δ H ð T 0 Þ ¼ H f ð C Þ þ H f ð D Þ − H f ð A Þ − H f ð B Þ : ð 16 : 26 Þ The terms H f ð C Þ etc. are the “ standard enthalpies of formation ” of each compound, for A, B, C, and D. A standard enthalpy of formation is the enthalpy change of the reaction needed to form 1 mole of A (or B, C, D, etc.) from the constituent molecules at their standard states at T 0 = 298 K and p = 1 atm. For example, H f for CO 2 is equal to H f ð CO 2 Þ ¼ H ð CO 2 Þ − H ð 1 mol of solid C Þ − H ð 1mol of gaseous O 2 Þ : ð 16 : 27 Þ Such values may be found in tables of data (see Appendix A , and Perry ’ s or the CRC Handbook). For molecules of a single component at its standard state, the value of H f is zero, by de fi nition.
eBook - PDF- Arther Adamson(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
Note that the phase (gas, liquid, solid) of the substance must be specified if there is any possible ambiguity. Unlike the situation with some extensive quantities, such as volume or mass, we have no absolute values for the internal energy or enthalpy of a substance. Although the first law permits us to calculate changes in Ε or H, its application never produces absolute values. The same is true in thermochemistry; ΔΕ and ΔΗ give the changes in internal energy and enthalpy that accompany a chemical reaction, and while either may be expressed in the form of Eq. (5-3), we do not know the separate Ε or 77 values. It is partly a consequence of this situation that much use is made of standard or reference states. Two systems of standard states are in use. The first is one of convenience, 1 atm pressure and, if so specified, 25°C. In the case of a reaction for which the reactants and products are all in a standard state, such as Eq. (5-4), one writes ΔΕ 29Β or ΔΗ 298 , where the superscript zero means that the pressure is one atm and the subscript gives the temperature chosen. As will be seen, a large body of thermochemical data is reported on this basis. The second system takes as the standard state the substance devoid of any thermal energy, that is, at 0 K. This is a more rational as well as a very useful approach and is developed in the Special Topics section. Its implementation does require either extensive knowledge of heat capacity data or sufficient spectroscopic information to allow evaluation of the various partition functions. 5-2 Measurement of Heats of Reaction: Relationship between ΔΕ and ΔΗ The practice of thermochemistry involves the measurement of the heat absorbed or evolved when a chemical reaction occurs. That is, the determination is one of q 5-2 MEASUREMENT OF HEATS OF REACTION; RELATIONSHIP BETWEEN ΔΕ AND ΔΗ 147 in the first law statements: dE = 8q — 8w, (5-6) dH = 8q + VdP.
eBook - PDF- R. Prasad(Author)
- 2016(Publication Date)
- Cambridge University Press(Publisher)
Thus change in the number of moles in the reaction D = –1 (ii) Conventionally all gases are treated as ideal gases, unless specified otherwise. The reaction has taken place at constant atmospheric pressure, denoted by P atm and at temperature T = 25 ° C which is equal to 298 K. For ideal gases, Application of Thermodynamics to Chemical Reactions 373 PV = RT or D( P atm V) = P atm D V = RD( T ) = RT ( D ) = RT (–1) S-9.2.1 While deriving the above equation we have made use of the fact that atmospheric pressure and temperature T remains constant in the reaction. We now apply Eq. 9.9, according to which, DQ P = DQ V + RD( T) S-9.2.2 In the above equation DQ P = DH, and DQ V = DU. What about the sign of DQ V ? It should be negative as heat is emitted in a combustion reaction. Hence, from Eq. S-9.2.2 we get, DH = – DU + (–1)RT = –13.68 × 105 J = –DU – (8.314 JK –1 ) × 298 K Or DU = 13.68 × 10 5 J – 2.48 × 10 3 J = 13.655 × 10 5 J 9.7 Relation between the Enthalpy Change in a Reaction and the Enthalpies of Formations of Reactants and Reaction Products It is easy to show that the enthalpy change in a reaction is equal to the difference between the enthalpies of formation of products and the reactants, i.e., DQ P = D D D H H H rea for for f f f = Â ( ) - Â ( ) products reactants 9.10 (a) Similarly, DQ V = D D D U U U = Â ( ) - Â ( ) products reactants 9.10 (b) It may further be remembered that measurements done using the bomb calorimeter are at constant volume, and, therefore, refers to the measurement of DQ V = DU Measurements done by a coffee calorimeter are done at constant pressure. It is, however, a common practice to specify the type of the measurement. Further, if the change in the standard enthalpy of a reaction is given it means it refers to constant pressure measurement. Another important point to remember is that enthalpy change in a reaction and in its inverse reaction is equal in magnitude but opposite in sign.
eBook - PDF- Arthur Adamson(Author)
- 2012(Publication Date)
- Academic Press(Publisher)
Note that the phase (gas, liquid, solid) of the substance must be specified if there is any possible ambiguity. Unlike the situation with some extensive quantities, such as volume or mass, we have no absolute values for the internal energy or enthalpy of a substance. Although the first law permits us to calculate changes in Ε or H, its application never produces absolute values. The same is true in thermochemistry; Δ Ε and ΔΗ give the changes in internal energy and enthalpy that accompany a chemical reaction, and while either may be expressed in the form of Eq. (5-3), we do not know the separate Ε or Η values. It is partly a consequence of this situation that much use is made of standard or reference states. Two systems of standard states are in use. The first is one of convenience, 1 atm pressure and, if so specified, 25°C. In the case of a reaction for which the reactants and products are all in a standard state, such as Eq. (5-4), one writes ΔΕ 298 or ΔΗ 298 , where the superscript zero means that the pressure is one atm and the subscript gives the temperature chosen. As will be seen, a large body of thermochemical data is reported on this basis. The second system takes as the standard state the substance devoid of any thermal energy, that is, at OK. This is a more rational as well as a very useful approach and is developed in the Special Topics section. Its implementation does require either extensive knowledge of heat capacity data or sufficient spectroscopic information to allow evaluation of the various partition functions. 5-2 Measurement of Heats of Reaction: Relationship between ΔΕ and AH The practice of thermochemistry involves the measurement of the heat absorbed or evolved when a chemical reaction occurs. That is, the determination is one of q
eBook - PDFChemistry
An Industry-Based Introduction with CD-ROM
- John Kenkel, Paul B. Kelter, David S. Hage(Authors)
- 2000(Publication Date)
- CRC Press(Publisher)
Just about all of the energy will be released as heat, though a little bit will also be done as work (see Section 13.2 for a review of “work”). For the remainder of this discussion, we will assume that our energy is heat energy. The term for heat energy at constant pressure (such as atmospheric pressure) is called enthalpy, H . We symbolize changes in enthalpy , which can be measured, by H. We will therefore say that the enthalpy change, or heat exchange, H, is 500 kJ for our reaction. We now see from the standpoint of energy exchange why this process can be used as an energy source. If we compare the energy exchange in the nuclear process to that in the chemical process, we see that changes in the nucleus provide vastly more energy per mole than chemical processes. This is why nuclear processes are so valued as power sources and in warfare, in spite of the attendant risks. Two final notes on determining energy exchange using bond energy values: 4moles of O–H bonds 467 kJ mole of O–H bonds 1868 kJ Chemistry Professionals at Work CPW Box 13.3 E NERGY M ANAGEMENT IN THE C HEMICAL I NDUSTRY he manufacture of chemicals through industrial chemical processes involves the manage-ment of energy going into the process as well as the management of energy coming from the process. Industrial companies consider energy an essential element of their economics—as much as, or more than other obvious factors such as land, labor, capital, and raw materials. Indeed, the proper management of energy dictates how the other factors are handled. The goal of the chemical process industry, like any other enterprise in the private sector, is to make money. Thus the cost of energy and quantity of energy required enter into the equation of a superior operation. Energy is often recycled, meaning that excess energy created in one part of the operation is often used in another part.
eBook - PDF- Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
- 2019(Publication Date)
- Openstax(Publisher)
INTRODUCTION CHAPTER 5 Thermochemistry 5.1 Energy Basics 5.2 Calorimetry 5.3 Enthalpy Chemical reactions, such as those that occur when you light a match, involve changes in energy as well as matter. Societies at all levels of development could not function without the energy released by chemical reactions. In 2012, about 85% of US energy consumption came from the combustion of petroleum products, coal, wood, and garbage. We use this energy to produce electricity (38%); to transport food, raw materials, manufactured goods, and people (27%); for industrial production (21%); and to heat and power our homes and businesses (10%). 1 While these combustion reactions help us meet our essential energy needs, they are also recognized by the majority of the scientific community as a major contributor to global climate change. Useful forms of energy are also available from a variety of chemical reactions other than combustion. For example, the energy produced by the batteries in a cell phone, car, or flashlight results from chemical reactions. This chapter introduces many of the basic ideas necessary to explore the relationships between chemical changes and energy, with a focus on thermal energy. Figure 5.1 Sliding a match head along a rough surface initiates a combustion reaction that produces energy in the form of heat and light. (credit: modification of work by Laszlo Ilyes) CHAPTER OUTLINE 1 US Energy Information Administration, Primary Energy Consumption by Source and Sector, 2012, http://www.eia.gov/totalenergy/ data/monthly/pdf/flow/css_2012_energy.pdf. Data derived from US Energy Information Administration, Monthly Energy Review (January 2014).
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