What does energetically stable mean

Therefore, the reaction would not occur without some outside influence such as persistent heating. However, endothermic reactions do occur spontaneously, or naturally. There must be another driving force besides enthalpy change which helps promote spontaneous chemical reaction. A very simple endothermic process is that of a melting ice cube.

Energy is transferred from the room to the ice cube, causing it to change from the solid to the liquid state. The solid state of water, ice, is highly ordered because its molecules are fixed in place. The melting process frees the water molecules from their hydrogen-bonded network and allows them a greater degree of movement. Water is more disordered than ice. The change from the solid to the liquid state of any substance corresponds to an increase in the disorder of the system.

There is a tendency in nature for systems to proceed toward a state of greater disorder or randomness. Entropy is a measure of the degree of randomness or disorder of a system. Entropy is an easy concept to understand when thinking about everyday situations. When the pieces of a jigsaw puzzle are dumped from the box, the pieces naturally hit the table in a very random state. In order to put the puzzle together, a great deal of work must be dome in order to overcome the natural entropy of the pieces.

The entropy of a room that has been recently cleaned and organized is low. As time goes by, it likely will become more disordered, and thus its entropy will increase see figure below. The natural tendency of a system is for its entropy to increase. Chemical reactions also tend to proceed in such a way as to increase the total entropy of the system. How can you tell if a certain reaction shows an increase or a decrease in entropy?

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. 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. Similarly, if you mix petrol gasoline and air at ordinary temperatures when you are filling up a car, for example , why doesn't it immediately convert into carbon dioxide and water?

It would be much more energetically stable if it turned into carbon dioxide and water - you can tell that, because lots of heat is given out when petrol burns in air. But there is no reaction when you mix the two. For any reaction to happen, bonds have to be broken, and new ones made. Breaking bonds takes energy. There is a minimum amount of energy needed before a reaction can start - activation energy. If the molecules don't, for example, hit each other with enough energy, then nothing happens.

We say that the mixture is kinetically stable , even though it may be energetically unstable with respect to its possible products. So a petrol and air mixture at ordinary temperatures doesn't react, even though a lot of energy would be released if the reaction took place. Petrol and air are energetically unstable with respect to carbon dioxide and water - they are much higher up the energy diagram.

But a petrol and air mixture is kinetically stable at ordinary temperatures, because the activation energy barrier is too high. If you expose the mixture to a flame or a spark, then you get a major fire or explosion. The initial flame supplies activation energy. The heat given out by the molecules that react first is more than enough to supply the activation energy for the next molecules to react - and so on.

The moral of all this is that you should be very careful using the word "stable" in chemistry! Note: You will find a bit more about activation energy on the introductory page about rates of reaction. If this is the first set of questions you have done, please read the introductory page before you start. Energy changes during chemical reactions Obviously, lots of chemical reactions give out energy as heat. Simple energy diagrams A reaction in which heat energy is given off is said to be exothermic.

A reaction in which heat energy is absorbed is said to be endothermic.