Which compound forms exothermically

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. This is clearly proportional to the activation energies of all possible reactions. As noted above, benzene is thermodynamically unstable compared with elemental carbon and hydrogen, but it is chemically stable under normal laboratory conditions, even when mixed with some reactive compounds such as bromine.

Compounds or mixtures that are chemically unstable are often called labile. When the rates of reactions and equilibria between reactants and products are carefully examined, it becomes apparent that overall enthalpy changes and enthalpy of activation barriers are not by themselves sufficient to explain the observations.

A modified energy function, called free energy , is needed. This thermodynamic function is described in the next two sections. Enthalpy is not the only thermodynamic function that influences the overall energy changes, rates and equilibria of chemical and physical transformations. Two examples will serve to demonstrate this fact.

First, nitrogen pentaoxide is an unstable solid that undergoes a spontaneous and endothermic decomposition to nitrogen dioxide and oxygen, as shown below.

Clearly, some factor other than a change in enthalpy must act to favor this decomposition. Second, the melting of solids e.

Since intermolecular attractions favor the liquefaction of gases and the solidification of liquids by lowering the enthalpy of the condensed phase, our world would be a frozen lump of solid matter were enthalpy the only controlling factor.

What do these two cases have in common that could account for their behavior? In the first example two well-defined molecules in the solid state break apart into five discrete product molecules, all of them in the gaseous state.

The second example likewise describes a progression from a highly ordered array of molecules in the solid state to a less ordered assemblage as a liquid, and finally to a nearly random disordered gaseous state. This disposition, favoring disorder, seems to be universal, as evidenced by the failure of a shuffled deck of cards to arrange itself into ordered suits. The thermodynamic name for this tendency toward randomness or disorder is entropy , symbol S. Entropy increases spontaneously, since greater randomness or disorder in a system has a higher statistical probability.

Returning to the deck of cards example, the number of different ways 52 cards may be arranged is very large 52! A disordered state is therefore more probable than an ordered one. Since the disorder created by molecular motion increases with temperature, the units of entropy, eu , are calories per degree Kelvin per mole. The standard state entropies of some elements and compounds are given in the following table.

Two obvious trends to note are that gases have higher entropies than liquids or solids of the same size, and molecules composed of many atoms have higher entropies than diatomic or triatomic molecules. Depending on the case to which they apply, these two functions may complement support each other or act in opposition. In careful studies of rates and equilibria the consequence of this relationship must be calculated. Examples of these will be found at the end of this chapter.

Ionization reactions in solution are complicated by solvation effects. The expected increase in entropy from the first factor listed above may be offset by solvent molecule orientation about the ions. The description of entropy given above reflects a traditional or classic view, which is presently undergoing a significant revision. In an exothermic reaction, energy is released because the total energy of the products is less than the total energy of the reactants. In the presence of water, a strong acid will dissociate quickly and release heat, so it is an exothermic reaction.

Endothermic reactions are reactions that require external energy, usually in the form of heat, for the reaction to proceed. Since endothermic reactions draw in heat from their surroundings, they tend to cause their environments to cool down.

They are also generally non-spontaneous, since endothermic reactions yield products that are higher in energy than the reactants. As such, the change in enthalpy for an endothermic reaction is always positive. In order to melt the ice cube, heat is required, so the process is endothermic.

Whether a reaction is endothermic or exothermic depends on the direction that it is going; some reactions are reversible, and when you revert the products back to reactants, the change in enthalpy is opposite. Boundless vets and curates high-quality, openly licensed content from around the Internet.

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Learning Objective Distinguish between endothermic and exothermic reactions.