Why is dissociation endothermic

For examples of direct and indirect determinations of heats of formation click on the above table. The first two equations show how the heats of formation of water and carbon dioxide are measured. Most heats of formation are negative, reflecting the strong covalent bonds and lower enthalpy that characterizes stable compounds relative to their elements. However, some stable compounds are found to have positive heats of formation, e. As we have noted, heats of reaction reflect the bond dissociation energies of bonds that are broken and formed in the reaction, but the formalism of setting elemental heats of formation to zero obscures the covalent bond dissociation energies of diatomic elements such as H 2 , O 2 , N 2 and Cl 2.

Elements that have solid standard states e. Fortunately, it is possible to determine the bond dissociation energy of diatomic elements and compounds with precision by non-thermodynamic methods, and together with thermodynamic data such information permits a table of average bond energies to be assembled. These bond energies or bond dissociation enthalpies are always positive, since they represent the endothermic homolysis of a covalent bond.

It must be emphasized that for the common covalent bonds found in polyatomic molecules e. C-H and C-C these are average dissociation enthalpies, in contrast to specific bond dissociation enthalpies determined for individual bonds in designated compounds. Factors such as hybridization, strain and conjugation may raise or lower these numbers substantially. Common sense suggests that molecules held together by strong covalent bonds will be more stable than molecules constructed from weaker bonds.

Previously we defined bond dissociation energy as the energy required to break a bond into neutral fragments radicals or atoms. The sum of all the bond energies of a molecule can therefore be considered its atomization energy , i.

If this concept is applied to the reactants and products of a reaction, it should be clear that a common atomization state exists, and that the total bond energies of the reactants compared with the bond energies of the products determines the enthalpy change of the reaction.

Thus, if the products have a greater total bond energy than the reactants the reaction will be exothermic, and the opposite is true for an endothermic reaction. The following diagram illustrates this relationship for the combustion of methane.

Always remember, a bond energy is energy that must be introduced to break a bond, and is not a component of a molecule's potential energy. Bond energies may be used for rough calculations of enthalpies of reaction. To do so the total bond energies of the reactant molecules must be subtracted from the total bond energies of the product molecules, and the resulting sign must be changed.

This operation is outlined above for the combustion of methane. To compare such a calculation with an experimental standard enthalpy of reaction, correction factors for heats of condensation or fusion must be added to achieve standard state conditions.

In the above example, gaseous water must be condensed to the liquid state, releasing Once this is done, a reasonably good estimate of the standard enthalpy change is obtained. It may be helpful to note that the potential energy of a given molecular system is inversely proportional to its total bond energies. In this reaction, potential energy is lost by conversion to kinetic heat energy. Thermodynamic calculations and arguments focus only on the initial and final states of a system.

The path by which a change takes place is not considered. Intuitively, one might expect strongly exothermic reactions to occur spontaneously, but this is usually not true. For example, the methane combustion described above does not proceed spontaneously, but requires an initiating spark or flame. Once begun, the heat produced by the combustion serves to maintain the reaction until one or both of the reactants are completely consumed.

Clearly, many potentially favorable reactions are prohibited or retarded by substantial energy barriers to the transformation. To understand why some reactions occur readily almost spontaneously , whereas other reactions are slow, even to the point of being unobservable, we need to consider the intermediate stages through which reacting molecules pass on the way to products. Every reaction in which bonds are broken will necessarily have a higher energy transition state on the reaction path that must be traversed before products can form.

This is true for both exothermic and endothermic reactions. In order for the reactants to reach this transition state, energy must be supplied from the surroundings and reactant molecules must orient themselves in a suitable fashion. Further treatment of this subject, and examples of reaction path profiles that illustrate transition states are provided elsewhere in this text.

However, in these introductory discussions a distinction between enthalpy and "potential energy" is not made. As expected, the rate at which chemical reactions proceed is, in large part, inversely proportional to their activation enthalpies, and is dependent on the concentrations of the reactants.

The study of reaction rates is called chemical kinetics. Common use of the term stability implies an object, system or situation that is likely to remain unchanged for a significant period of time.

In chemistry, however, we often refer to two kinds of stability. Thermodynamic Stability : The enthalpy or potential energy of a compound relative to a reference state. For exothermic reactions we may say that the products are thermodynamically more stable than the reactants.

The opposite would be true for endothermic reactions. Chemical Stability : The resistance of a compound or mixture of compounds to chemical change reaction.

Access now. Your report is submitted, the author will be informed about it. Question Followers. Report Question. Can be made better Question lacks the basic details making it difficult to answer. Spam Question drives traffic to external sites for promotional or commercial purposes.

Irrelevant The Question is not relevant to User. Report Answer. Spam Answer drives traffic to external sites for promotional or commercial purposes.

Does not Answer the Question Answer posted is not solving the query properly. Poorly Written Answer is not written properly. Factually Incorrect Facts stated in the Answer are incorrect. Edit Comment. Delete Comment.

Toward a general theory of heterogeneous reactions: thermochemical approach. Interpretation of the kinetic compensation effect in heterogeneous reactions: thermochemical approach.

Int Rev Phys Chem. Galwey AK. Theory of solid-state thermal decomposition reactions scientific stagnation or chemical catastrophe? An alternative approach appraised and advocated. Solid state reaction kinetics, mechanisms and catalysis: a retrospective rational review. React Kinet Mech Catal. Evaluation of thermodynamic functions from equilibrium constants.

Trans Faraday Soc. Measurement and correlation of the solubility of enrofloxacin in different solvents from J Chem Eng Data. New static apparatus and vapor pressure of reference materials: naphthalene, benzoic acid, benzophenone, and ferrocene. Fluorene: an extended experimental thermodynamic study. J Chem Thermodyn. Thermal properties of amine hydrochlorides. Part I. Thermolysis of primary n -alkylammonium chlorides.

Thermal features thermolysis and thermochemistry of hexachlorostannates of some mononitrogen aromatic bases. Mianowski A, Siudyga T. RSC Adv. De Marco D, Linert W. Thermodynamic relationships in complex formation. Ryde U. A fundamental view of enthalpy—entropy compensation.

Med Chem Commun. The thermodynamics of protein—ligand interaction and solvation: insights for ligand design. J Mol Biol. Klebe G. Applying thermodynamic profiling in lead finding and optimization. Nat Rev Drug Discov. J Appl Sol Chem Model. International Union of Pure and Applied Chemistry. Compendium of chemical terminology. Gold Book. Version 2. Accessed 15 Oct CAS Google Scholar.

Mianowski A. Thermal dissociation in dynamic conditions by modeling thermogravimetric curves using the logarithm conversion degree. Analysis of the thermokinetics under dynamic conditions by relative rate of thermal decomposition.

J Therm Anal Cal. Some remarks on equilibrium state in dynamic condition. Barin I. Thermochemical data of pure substance. Weinheim: VCH; Google Scholar. Odak B, Kosor T. Quick method for determination of equilibrium temperature. Computational prediction of the pattern of thermal gravimetry data for the thermal decomposition of calcium oxalate monohydrate. Download references. We would like to express our sincere gratitude to Professor Evgeni B.

Starikov for his attentive and critical reading of our final version of manuscript and for his substantive and editorial support. You can also search for this author in PubMed Google Scholar. Reprints and Permissions. Mianowski, A. Thermal dissociation in terms of the second law of chemical thermodynamics. J Therm Anal Calorim , — Download citation. Received : 05 November Accepted : 14 May Published : 18 June Issue Date : November