Chemistry
Bond Enthalpy
Bond enthalpy is the energy required to break a chemical bond. It is a measure of the strength of a bond and is typically expressed in kilojoules per mole (kJ/mol). When a bond is broken, energy is absorbed, and when a bond is formed, energy is released. Bond enthalpy values are useful for predicting the energy changes in chemical reactions.
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eBook - ePub- 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;ΔH b= Σ ΔH bd/ number of bonds =120 kcal •/ 2 bondsmol+ 101 kcal •− 1mol− 1= 110.5 kcal •mol− 1Table 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
eBook - ePub- 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
eBook - ePub- Adrian Dingle, Derrick C. Wood(Authors)
- 2014(Publication Date)
- Research & Education Association(Publisher)
6. Making bonds is an exothermic process that releases energy (exothermic changes are given negative values).7. Breaking or making a particular bond involves the same magnitude of energy change, i.e., the energy change is the same for breaking and making a particular bond, all that differs is the sign. For example, an average bond energy of 347 kJ for the C–C bond, means that 347 kJ of energy will be released (–347) when the bond is broken in a reaction, and that 347 kJ of energy is required to break the bond (+347) in a reaction.
A. Breaking and Making Bonds 1. All chemical reactions involve the making and breaking of bonds. 2. The energy required to break all of the reactant bonds is a sum of all of the average bond energies in the reactants. This is an endothermic term. 3. The energy released in making all of the product bonds is a sum of all of the average bond energies in the products. This is an exothermic term.II. Bond Energies in Chemical Reactions4. The energy change in the reaction as a whole is the sum of the breaking (endothermic) process and the making (exothermic) process, and can be negative or positive depending on the relative magnitude of the two (breaking and making) processes.5. In an exothermic reaction, thermal energy is transferred to the surroundings. In an endothermic reaction, thermal energy is transferred from the surroundings. In each case the energy transfer is called the enthalpy change or ΔH, and it is usually measured in kJ/mol.System refers to a particular part of the universe that is being studied, and, in chemistry, this usually simply means a certain, chemical reaction.Surroundings is everything that is outside of the system.Universe is the system AND the surroundings.Endothermic is energy transferred from the surroundings to the system.Exothermic is energy transferred from the system to the surroundings.Standard Enthalpies of Formation is the energy change when 1 mole of a substance is formed from its elements in their standard states. Also called ΔH°f
- 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.
SOLUTION:Bond Enthalpy(kJ/mol) H—H 436 O=O 495 H—O 464 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 YOUKNOW?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.
eBook - ePub- 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.
