Bond Energy Calculations

5 Key excerpts on "Bond Energy Calculations"

• 

  eBook - ePub

  General Chemistry for Engineers

  • Jeffrey Gaffney, Nancy Marley(Authors)
  • 2017(Publication Date)
  • Elsevier

    (Publisher)

  bond dissociation energy , is the standard energy required to break one specific bond in a molecule in the gas phase. The bond dissociation energy of a specific chemical bond in a molecule depends on the molecular environment surrounding the bond. The bond energy values usually given for a particular kind of chemical bond are values averaged over different environments.

  Water is a good example of how these two related energy terms differ from one another. The bond dissociation energy of the first O H bond is 120 kcal • mol− 1 at 298 K. This value is the bond dissociation energy (ΔH °bd ) for the HO H bond. But, the bond dissociation energy for the second O H bond is 101 kcal • mol− 1 at 298 K. This value is the bond dissociation energy for the O H bond. Thus, the bond energy (ΔH °b ) for the O H bonds in the water molecule is the average of the bond dissociation energies for the two O H bonds. So, the bond energy for water is the sum of the two ΔH bd values for the O H bonds divided by the number of bonds, or;

  Δ

  H b

  = Σ Δ

  H bd

  / number of bonds =

  120 kcal •

  mol

  − 1

  + 101 kcal •

  mol

  − 1

  / 2 bonds

  = 110.5 kcal •

  mol

  − 1

  Table 8.3 gives some bond energies for selected types of chemical bonds. Bond energies are always positive as it requires energy to break a bond. The sign of the energy required to form a bond is always negative because energy is released when bonds are formed. The energy required to form a specific bond is the negative value of the bond energy for the same bond. The change in enthalpy during a chemical reaction (ΔH rxn

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  eBook - ePub

  Philosophy of Chemistry

  • Dov M. Gabbay, Paul Thagard, John Woods, Dov M. Gabbay, Paul Thagard, John Woods(Authors)
  • 2011(Publication Date)
  • North Holland

    (Publisher)

  The energetic conception of the bond is the logical outcome of Coulson's sceptical thoughts about the bond, and the breakdown of valence formulae in many compounds. Rather than seeking a material part that realises the theoretical role of keeping a molecule together, the energetic conception fixes on what is common to all cases of chemical bonding: changes in energy. On the energetic view, facts about chemical bonding are just facts about energy changes between molecular or super-molecular states. There is no requirement, or motivation, for bonds to be localized or localizable within the molecule, or directional. Hence the energetic view is more general and agnostic than the structural view, and is more a theory of chemical bonding than a theory of bonds.

  The energetic view finds support outside quantum mechanics. In thermodynamics, the strength of bonds can be estimated using Hess' law, according to which the change in enthalpy between two states is independent of the specific path taken between them. A measure of the strength of the carbon-hydrogen bond in methane, for instance, can be estimated by breaking the process of the formation of methane down into formal steps: the atomisation of graphite and molecular hydrogen followed by the formation of methane. The heat of formation of methane from graphite and molecular hydrogen, and also the heats of atomisation of graphite and molecular hydrogen are empirically measurable, and the energy change in the formation of the C-H bonds is just the difference between them. In similar fashion, the lattice energy of common salt is the change in enthalpy when the salt lattice is formed from the gaseous ions Na+ and Cl− .

  Within quantum mechanics, the molecular orbital approach represents the formation of a bond by correlating the electronic configuration of the molecule's bonded state with that of the separated atoms. A bond is formed just in case the bonded state is of lower energy than the separated atoms. The electronic configuration is described in terms of delocalised molecular orbitals, spread over the entire molecule, and the molecular geometry is explained as a local minimum on a potential energy surface, determined by the dependence on geometry of the occupied orbital energies.

  The energetic view says no more than this about bonds. While it has the advantage that it applies straightforwardly to all types of bonding, and is clearly consistent with quantum mechanics, the view dispenses with much of what is plausible and intuitive in the structural conception of the bond. Localised, directed bonds have been central to understanding the structure, symmetry properties, spectra and reactivity of organic molecules since the nineteenth century. They remain central to modern organic chemists' understanding of reaction mechanisms. The energetic conception presents a radically revised view of the bond, and it remains to be shown that it can provide alternative explanations that do not appeal to the classical bond. If it cannot, the energetic conception implies explanatory loss.

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  eBook - ePub

  AP® Chemistry All Access Book + Online + Mobile

  • Kevin Reel, Derrick C. Wood, Scott A. Best(Authors)
  • 2014(Publication Date)
  • Research & Education Association

    (Publisher)

  exothermic process. Heats of reaction can be estimated by finding the difference between the energy required to break all of the bonds in a molecule and the energy released when the bonds are formed. You will first have to draw the molecular structures of the reactants and products, and then evaluate the types of bonds broken and formed during a reaction.

  EXAMPLE: Use the following table of bond energies to approximate the change in enthalpy when 1 mole of hydrogen is combusted.

  Bond	Enthalpy(kJ/mol)
  H—H	436
  O=O	495
  H—O	464

  SOLUTION:

  There are (2) H—H bonds and (1) O = O bond to break in the reactants.

  There are a total of (4) O—H bonds formed in the products.

  TEST TIP For complex reactions, you can perform Bond Energy Calculations just based on the bonds that are actually broken and the bonds that are formed. This will save you some precious time.

  Energy and Phase Changes

  When performing energy calculations, the state of matter of the reactants and products is highly important. Particles in the solid state have very little energy compared to those in the gaseous state. In addition, when a substance changes phase, there is an energy associated with those transitions. For the solid–liquid phase change, the enthalpy of fusion (Hfus ) is used to quantify the amount of heat required to melt a solid. This value also tells you how much energy is released when a particle freezes. Similarly, the enthalpy of vaporization (Hvap ) describes the energy required for the liquid-gas phase change. This value also tells you the amount of energy released when a gas is condensed to the liquid state.

  DID YOU

  KNOW?

  Disposable hand warmers are created from solid iron that is spread over an enormous surface area. When the iron-containing hand warmer is exposed to air, it undergoes a redox reaction that releases a lot of heat and forms iron (III) oxide, which is commonly known as rust.

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  eBook - ePub

  Foundations for Teaching Chemistry

  Chemical Knowledge for Teaching

  • Keith S. Taber(Author)
  • 2019(Publication Date)
  • Routledge

    (Publisher)

  10 Energy in chemistry and chemical bonding

  This chapter discusses one of the key topics in the chemistry curriculum, chemical bonding. This is a highly abstract concept area where a range of models and simplifications are taught. It is also an area where students commonly develop tenacious alternative conceptions (Taber, 2013a), and thus where the teaching approach can be very important in channelling student thinking towards scientific models. One particular feature of many students’ thinking is that they learn about chemistry topics such as bonding with no cognisance of the basic physical principles they have been taught elsewhere in science. Yet if students are to develop scientific understandings of chemistry, they need to appreciate where key concepts from physics, such as force and energy, are applied. This chapter reflects this imperative by first considering the role of the energy concept in understanding school chemistry before specifically addressing chemical bonding concepts.

  Appreciating the physicists’ concept of energy and how this applies in chemistry

  Energy is one of the most fundamental and ubiquitous concepts in science. It is also one of the most abstract. It is closely associated with another abstract concept – force. The primary responsibility for teaching these ideas falls upon the physics teacher. There are, however, consequences here for the teacher of chemistry:

  1. To avoid the potential of confusing students, the science department as a whole should have a common way of talking about energy and force so the ideas are used consistently across different topics and science subjects.
  2. As energy is an important concept in chemistry, the teacher of chemistry has to rely on what has been taught and how it has been taught in another subject.
  3. The teacher of chemistry not only relies on what has been taught in physics but on whether students can transfer their learning in physics to other subjects.

  The latter consideration is not insignificant. The question of transfer of learning, applying what has been learnt beyond the original context, is seen as a major issue in education (Lobato, 2006) because students often struggle to apply what they have learned in one context in other relevant situations. This is an example of a ‘fragmentation’ learning impediment (see Figure 3.1 ), as the learner does not make intended links with prior learning (see Figure 10.1

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  eBook - ePub

  Understanding General Chemistry

  • Atef Korchef(Author)
  • 2022(Publication Date)
  • CRC Press

    (Publisher)

  It is also a measure of the bond strength. Bond enthalpy values are always positive since bond breaking is an endothermic process. Example: The bond enthalpy of the H–H bond, ΔH(H–H), is equal to 436 kJ mol −1, which means that we need an energy of 436 kJ to break the H–H bonds in one mole of H–H bonds. A chemical reaction can be described as the breaking of bonds in the reactants and the making of bonds in the products. Thus, we can calculate the standard enthalpy of the reaction, Δ H rxn ° : Δ H r x n ° = ∑ Δ H b r o k e n b o n d s − ∑ Δ H f o r m e d b o n d s Example: For the ethene hydrogenation. reaction: C 2 H 4 g + H 2 g → C 2 H 6 g, Δ H rxn ° = ? Δ H rxn ° = Δ H C=C + 4 × Δ H C − H + Δ H H − H − Δ H C − C + 6 × Δ H C − H Lattice energy can be defined in two opposite ways: It is the amount of energy that is spent to separate an ionic crystal into its constituent gaseous ions. It is the energy released when gaseous ions bind to form an ionic compound. This process is exothermic, and the values for lattice energy, expressed in kJ mol −1, are negative. The lattice energy can be calculated by applying Hess’s law on a series of individual reactions, forming a cycle. This cycle is called the Born–Haber cycle. The Born–Haber cycle is used to calculate the lattice enthalpy of an ionic compound, formed by a metal and a non-metal. It can be obtained by the following steps: Sublimation of the metal Dissociation of the gaseous non-metal Ionization of the gaseous metal Formation of the gaseous non-metal anion Formation of the ionic compound Example: The Born–Haber cycle for NaCl Δ H 1 ° = Δ H sub Na + 1 2 × Δ H diss Cl 2 + I Na + A e Cl + Δ H Lattice Δ H Lattice = Δ H 1 ° − Δ H sub Na + 1 2 Δ H diss Cl 2 + I Na + A e Cl = − 788 kJ

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