Enthalpy of Formation

11 Key excerpts on "Enthalpy of Formation"

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  eBook - ePub

  Handbook of Thermal Analysis and Calorimetry

  Recent Advances, Techniques and Applications

  • (Author)
  • 2011(Publication Date)
  • Elsevier Science

    (Publisher)

  Chapter 14

  Thermochemistry

  M.V. Roux and M. Temprado,     Institute of Physical-Chemistry “Rocasolano”, C.S.I.C., Serrano 119, 28006-Madrid, Spain

  1 Introduction

  1.1 The objectives of thermochemistry

  Two basic properties of chemical compounds are the structures and the energies of their molecules. These are intimately related because the energy associated with a particular structure depends on the atoms, types of bonds and angles that form the structure. The thermochemist is interested in enthalpy changes accompanying reaction, but even more in the enthalpies of formation of compounds from their elements. This valuable and fundamental thermodynamic property of a material is defined as the enthalpy change that occurs upon the formation of a compound from its component elements in their standard states, at a determined temperature of reference, usually 298.15 K, and a standard pressure, currently taken to be 101.325 kPa (1 atmosphere). Values of enthalpies of formation provide a measure of the relative thermochemical stabilities of molecules, intimately related to their structures. It is, however, necessary to eliminate the intermolecular and network energies and to refer the enthalpies of formation to the gaseous state. From values of the enthalpies of formation of the molecules in gaseous state, the steric, electronic and electrostatic effects of different substituents can be evaluated. Comparison of the enthalpies of formation of isomeric compounds is particularly useful because it shows their relative stabilities and provides evidence on the interactions that are responsible for the Enthalpy of Formation of each isomer. For example, Figure 1 depicts graphically the enthalpies of formation of the three isomers of formula C5 H12 , pentane [1 ], 2-methylbutane [1 ] and 2,2-dimethylpropane [1 ]. All of them are more stable (have lower enthalpy) than five carbon atoms and six hydrogen molecules in their standard states. Nevertheless, the enthalpies of formation of the pentane isomers reveal that 2-methylbutane is more stable than pentane by 6.9 kJ mol−1 and the most branched hydrocarbon, 2,2-dimethylpropane, is 21.1 kJ mol−1

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  General Chemistry: Atoms First

  • Young, William Vining, Roberta Day, Beatrice Botch(Authors)
  • 2017(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  c 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 297 10.6 Standard Heats of Reaction 10.6a Standard Heat of Formation The enthalpy change for a chemical reaction varies with experimental conditions such as temperature, pressure, and solution concentration. Therefore, it is helpful to specify a standard set of conditions under which to tabulate enthalpy change, as well as a standard set of enthalpy change values that can be used to calculate the enthalpy change for a reaction. Enthalpy change values that are tabulated under standard conditions are indicated with a superscripted ° symbol, as in D H ° rxn . Standard state conditions for enthalpy values are: ● Gases, liquids, and solids in their pure form at a pressure of 1 bar at a specified temperature ● Solutions with concentrations of 1 mol/L at a specified temperature Notice that standard state conditions do not specify a temperature. Most standard enthalpy values are tabulated at 25 °C ( 298 K), but it is possible to report standard enthalpy change values at other temperatures. Creating a list of all possible standard enthalpy values is impossible. Fortunately, tabulating only one type of standard enthalpy change is all that is needed to calculate the standard enthalpy change for almost any chemical reaction. The standard heat of formation (or standard Enthalpy of Formation) for a species is the enthalpy In order to use Hess’s law, the sum of two or more chemical reactions must result in the net reac-tion. In this example, reversing the second reaction and adding it to the first reaction results in the net reaction. Notice that reversing the second reaction requires changing the sign of D H (2). Because the net reaction can be expressed as the sum of two chemical reactions, the enthalpy change for the net reaction is equal to the sum of the enthalpy changes for the individual steps.

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  A Textbook of Physical Chemistry

  • Arthur Adamson(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  ^ # 2 9 8 = -337.23 kcal, (5-18) ^ # 2 ° 9 8 = -68.317 kcal, (5-19) ^ # 2 ° 9 8 = -372.82 kcal. (5-20) 5-5 ENTHALPIES OF FORMATION 153 5-5 Enthalpies of Formation A. Standard Enthalpy of Formation for Compounds A particularly useful way of summarizing thermochemical data is in terms of the standard Enthalpy of Formation. The reaction is that of formation of the compound from the elements, each element being in its stable chemical state and phase at 25°C and 1 atm pressure. For example, the reaction giving the standard Enthalpy of Formation of H 2 0(/), ΔΗ^ 298 ( H 2 0 , /), is H 2 (g, 298 Κ, 1 atm) + hO z (g, 298 Κ, 1 atm) = H 2 0(/, 298 Κ, 1 atm). (This particular enthalpy change is also that of the combustion of hydrogen.) That for methane is C(graphite, 298 Κ, 1 atm) + 2U 2 (g, 298 Κ, 1 atm) = CH^g, 298 Κ, 1 atm), (5-22) The heat of formation of an element in its standard state is zero by definition. It is not always possible to determine a heat of formation directly, as in the case of water, and the values are usually calculated by the indirect procedure. Thus J / / ° 2 9 8 ( C H 4 ) could be obtained by application of Eq. (5-21) to the heats of combustion of graphite, hydrogen, and methane. Extensive tables of standard enthalpies of formation have been built up, and a sampling is given in Table 5-2. The table includes values for unstable forms of some elements, such as 0 3 , and C (diamond) and in some cases also gives values for both gaseous and liquid states, as for water. The difference between two such values is simply the enthalpy of vaporization. One may calculate the standard enthalpy change for any chemical reaction involving substances whose standard enthalpies of formation are known.

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  Classical and Quantum Thermal Physics

  • R. Prasad(Author)
  • 2016(Publication Date)
  • Cambridge University Press

    (Publisher)

  In case the substance is in the solid or in the liquid phase the standard state is the pure form of the substance at the atmospheric pressure and at 298 K (= 25 ° C) temperature. If the substance is a pure gas then the standard state is at one atmospheric pressure and at temperature of 298 K, unless specified otherwise. If there is a mixture of gases than the standard state of each gas is at a partial pressure of 1 atmosphere and the specified temperature which is generally 298 K. For a substance in solution the standard state refers to 1-molar concentration. 9.5 The Standard Enthalpy Change, DH rea Δ , in a Chemical Reaction In a chemical reaction, on the left hand side of the reaction equation, are reactants and on the right hand side the reaction products. Reactants Æ Reaction products 372 Classical and Quantum Thermal Physics The standard enthalpy change in such reactions refer to the change in enthalpy when specified number of moles of the reactants, all under standard states, are converted completely to the specified number of moles of products, all at standard states. The standard enthalpy change for a reaction is generally denoted by DH rea Δ , here superscript f denots that it refers to standard conditions of the reactants and the products and the subscript ‘rea’ denotes that it is for the specified reaction. 9.6 Standard Molar Enthalpy of Formation (Heat of Formation) The standard molar enthalpy DH for f of a substance is the change in enthalpy in the reaction in which one mole of the substance in specified state is formed from its elements in standard states. By convention the standard molar enthalpy of elements in their standard states is taken to be zero. For example, it is known that 1 2 mole of H 2 in gaseous state plus 1 2 mole of Br 2 in liquid state on reaction produce 1 mole of HBr 2 in gaseous state.

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  A Textbook of Physical Chemistry

  • Arther Adamson(Author)
  • 2012(Publication Date)
  • Academic Press

    (Publisher)

  The reaction is that of formation of the compound from the elements, each element being in its stable chemical state and phase at 25°C and 1 atm pressure. For example, the reaction giving the standard Enthalpy of Formation of H 2 O(0, ΔΗ^ 29% ( H 2 0 , /), is H 2 (g, 298 Κ, 1 atm) + iO z (g 9 298 Κ, 1 atm) = H 2 0 ( / , 298 Κ, 1 atm). (This particular enthalpy change is also that of the combustion of hydrogen.) That for methane is Qgraphite, 298 Κ, 1 atm) + 2 H 2 ( ^ , 298 Κ, 1 atm) = C H 4 ( ^ , 298 Κ, 1 atm), (5-22) The heat of formation of an element in its standard state is zero by definition. It is not always possible to determine a heat of formation directly, as in the case of water, and the values are usually calculated by the indirect procedure. Thus ^7/f °,298 (CH 4 ) could be obtained by application of Eq. (5-21) to the heats of combustion of graphite, hydrogen, and methane. Extensive tables of standard enthalpies of formation have been built up, and a sampling is given in Table 5-2. The table includes values for unstable forms of some elements, such as 0 3 , and C (diamond) and in some cases also gives values for both gaseous and liquid states, as for water. The difference between two such values is simply the enthalpy of vaporization. One may calculate the standard enthalpy change for any chemical reaction involving substances whose standard enthalpies of formation are known. The The example may be generalized. We may obtain ΔΗ 29Β f ° r a n v reaction by adding the standard enthalpies of combustion of tne reactants and subtracting the sum of those of the products: ΔΗ°2*8= Σ Γ ^ 2 , 2 9 8 -Σ P ^ ° C 2 9 8 , (5-21) reactants products where ΔΗ^ 29Β denotes a standard enthalpy of combustion. As an illustration, AH 29B for the (unlikely) reaction 2CH 4 (£) + CH 3 COOH(/) = C 3 H 8 t e ) + C0 2 (g) + 2U 2 (g) is given by 2 AH° Ct29S (CH 4 ) + J / / C ° , 2 9 8 ( C H 3 C O O H ) -AH° Ct29S (C 3 H 8 ) - 2 AH° Ct29S (H 2 ).

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  Chemistry, 5th Edition

  • Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
  • 2022(Publication Date)
  • Wiley

    (Publisher)

  Figure 8.16 shows how we can use the standard Enthalpy of Formation, Δ f H o , to calculate the atomisation enthalpy. Across the bottom we have the chemical equation for the formation of 1 mol of CH 4 from its elements in their standard states. The enthalpy change for this reaction is Δ f H o CH 4 (g) . In this figure, we also can see an alternative three-step path that leads to CH 4 (g). One step is the breaking of HH bonds in the H 2 molecules to give 4 mol of gaseous hydrogen atoms; another is the vaporisation of carbon to give 1 mol of gaseous carbon atoms, and the third is the combination of the gaseous atoms to form 1 mol of gaseous CH 4 molecules. These changes are labelled 1, 2 and 3 in the figure. FIGURE 8.16 Two paths (equations) for the formation of methane from its elements in their standard states. As described in the text, steps 1, 2 and 3 of the upper path are the formation of gaseous atoms of the elements and then the formation of the bonds in CH 4 . The lower path is the direct combination of the elements in their standard states to give CH 4 . Because ΔH is a state function, the sum of the enthalpy changes along the upper path must equal the enthalpy change for the lower path, which is Δ f H o . 4H(g) C(g) C(s) + + 2H 2 (g) CH 4 (g) 3 1 2 Since enthalpy is a state function, the net enthalpy change from one state to another is the same regardless of the path that we follow. This means that the sum of the enthalpy changes along the upper path must be the same as the enthalpy change along the lower path, which is Δ f H o . Perhaps this can be seen more easily in Hess’s law terms if we write the changes along the upper path in the form of thermochemical equations. Steps 1 and 2 have enthalpy changes that are standard enthalpies of formation of gaseous atoms. Values for these quantities have been measured for many of the elements, and some are given in table 8.4.

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  Physical Chemistry

  Thermodynamics

  • Horia Metiu(Author)
  • 2006(Publication Date)
  • Taylor & Francis

    (Publisher)

  For example, the heat of the reaction N(g) + 3H(g) → NH 3 (g) is not the heat of formation of NH 3 , because under standard conditions nitrogen and hydrogen are diatomics, not atoms. The heat of the reaction Zn(g) + 1 2 O 2 (g) → ZnO(s), in which the Zn vapor reacts with oxygen, is not the heat of formation of solid zinc oxide (ZnO), because Zn is solid under standard conditions. Occasionally, there are exceptions to these rules. For example, you will find in tables the heat of formation of CH 3 OH(g) even though CH 3 OH is liquid under standard conditions. It is also likely to find the in the tables the heat of formation of CO 2 (g) from either graphite or diamond. In such cases, the author of the table will alert you that liquid CH 3 OH or diamond has been used in the reaction for which the heat of formation is given. According to our definition, the heat of formation of an element in the state of aggregation that is stable under standard conditions is zero. For example, under standard conditions the stable form of carbon is graphite, not diamond. The heat of formation of graphite is 0 and that of diamond is 1.895 kJ/mol. §26. A Bit of Preparation. To prepare for the proof connecting the heat of a reaction to the heats of formation of the participating compounds, I will examine a very simple “reaction,” A → C, whose heat of reaction is H . You happen to know that the heat of the reaction A → B is H 1 , and that of B → C is H 2 . If you look at Fig. 13.1, you realize that you can perform the reaction A → C on two paths. One is direct and has the heat H ; the other is indirect (A → B → C) and has the heat H 1 + H 2 , but the change of enthalpy is independent of path so H = H 1 + H 2 . You can also prove this equation by using the definitions H = h (C) − h (A), H 1 = h (B) − h (A), and H 2 = h (C) − h (B). If you add H 1 + H 2 = ( h (B) − h (A) ) + ( h (C) − h (B) ) , you get H = h (C) − h (A).

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  General Chemistry I as a Second Language

  Mastering the Fundamental Skills

  • David R. Klein(Author)
  • 2015(Publication Date)
  • Wiley

    (Publisher)

  Let’s say you look up in a chart that the H for a particular reaction is 50 kJ/mol. If you are run- ning the reaction on 2 moles, then the H will be 100 kJ. If you use the value 50 kJ in your calculations, then you will get the wrong answer. Notice that the units are kJ/mol. The /mol is an indication that this property is extensive. Now that we have seen what enthalpy is and how it is measured, let’s take a look at some common terminology that chemists use when referring to enthalpy. We have already argued that, under conditions of constant pressure, H  q. Re- member that q refers to the transfer of energy in the form of heat. Therefore, we often refer to H as the “heat of reaction”. That is a common term used for H, so you should get used to it now. In general, the H for any reaction is called the heat of reaction, but there are a few specific categories of reactions that get their own special name for H. Here are the two common names that you will see in this course: 1. When one mole of a substance reacts with O 2 , the process is a called a com- bustion reaction. The change in enthalpy associated with a combustion reac- tion is called the heat of combustion, and it is shown like this: H comb . 2. When one mole of a substance is formed from its elements, the change in en- thalpy associated with the reaction is called the heat of formation, and it is shown like this: H f . The two examples above are special types of reactions. But there are also a couple of special processes (not reactions) that also get special names. Here are the two common names that you will see in this course: 1. When one mole of a substance turns from a solid into a liquid (so the solid is melting), there is a change in enthalpy associated with the melting process. This H is called the heat of fusion, and it is shown like this: H fus .

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  Chemistry for Technologists

  The Commonwealth and International Library: Electrical Engineering Division

  • G. R. Palin, N. Hiller(Authors)
  • 2014(Publication Date)
  • Pergamon

    (Publisher)

  Heat Content or Enthalpy (H) A gas has another form of energy in addition to those comprising the internal energy. This is the ability to do work by expansion, and depends on its state of compression. A given amount of a gas at a fixed tempera-ture has a fixed internal energy, but the higher its pressure, the greater is its potential to do work by expanding. In order to allow for this type of energy it is necessary to define another function, known as the heat con-tent, or the enthalpy, denoted by H, and such that Η = U + PV As with internal energy, when a change occurs in a system, the change in enthalpy depends only on the initial and final states, and AH = //final — ^initial If a change occurs in a system and there is no change in pressure and volume, the changes in internal energy and enthalpy are the same. If the pressure remains constant, but the volume changes, AH = AU + PAV 62 C H E M I S T R Y F O R T E C H N O L O G I S T S Heat of Reaction Consider an exothermic reaction occurring at constant volume, and under conditions such that there is no heat lost from the system. The final state is the products of the reaction at a higher temperature than the initial reactants. The chemical energy of the reaction has been converted into thermal energy, and absorbed by the products. The useful energy of the reaction can be considered as the amount of heat which would have to be removed from the system to maintain the temperature constant. This is equal in magnitude, and opposite in sign, to the change in internal energy of the system. Useful energy = — Δ U (at constant temperature). This relationship also holds for an endothermic reaction, but in this case the useful energy is the amount of heat which must be introduced to bring about the reaction at constant temperature and volume.

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  Chemistry

  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 .

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  Physical Chemistry

  Understanding our Chemical World

  • Paul M. S. Monk(Author)
  • 2005(Publication Date)
  • Wiley

    (Publisher)

  And, as we have consistently reported, the best macroscopic indicator of a microscopic energy change is a change in temperature. Like internal energy, we can never know the enthalpy of a reagent; only the change in enthalpy during a reaction or process is knowable. Nevertheless, we can think of changes in H . Consider the preparation of ammonia: N 2(g) + 3H 2(g) −−→ 2NH 3(g) (3.26) We obtain the standard enthalpy change on reaction H O r as a sum of the molar enthalpies of each chemical participating in the reaction: H O r = products νH O m − reactants νH O m (3.27) The values of ν for the reaction in Equation (3.26) are ν (NH 3 ) = +2, ν (H 2 ) = −3 and ν (N 2 ) = −1. We obtain the standard molar enthalpy of forming ammonia after inserting values into Equation (3.27), as H O r = 2H O m (NH 3 ) − [H O m (N 2 ) + 3H O m (H 2 )] SAQ 3.8 Write out an expression for H O r for the reaction 2NO + O 2 → 2NO 2 in the style of Equation (3.27). Unfortunately, we do not know the enthalpies of any reagent. All we can know is a change in enthalpy for a reaction or process. But what is the magnitude of this energy change? As a consequence of Hess’s law (see p. 98), the overall change in enthalpy accompanying a reaction follows from the number and nature of the bonds involved. We call the overall enthalpy change during a reaction the ‘reaction enthalpy’ H r , and define it as ‘the change in energy occurring when 1 mol of reaction occurs’. In consequence, its units are J mol −1 , although chemists will usually want to express H in kJ mol −1 . In practice, we generally prefer to tighten the definition of H r above, and look at reagents in their standard states. Furthermore, we maintain the temperature T at 298 K, and the pressure p at p O . We call these conditions standard temperature and pressure, or s.t.p. for short. We need to specify the conditions because temperature and pressure can so readily change the physical conditions of the reactants and products.

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