This bond may be broken by the introduction of heat or light energy, and it has been determined that This is shown in the first of the following reactions; the second reaction describes the reverse bond-forming process. The energy absorbed or released in these reactions is referred to as the bond dissociation energy.
If the bond dissociation energy is introduced from the surroundings in the form of heat, the transformation is said to be endothermic. If heat passes from the system to the surroundings, the transformation is termed exothermic. Using our initial terminology, we may say that the covalently bonded system has a lower potential energy than the unbonded diatomic system. Indeed, It is helpful to think of exothermic reactions as proceeding from a higher energy less stable reactant state to a lower energy more stable product state, as shown in the diagram on the right.
Some basic principles of reaction energetics were discussed earlier. In more complex chemical reactions some or even all of the bonds that hold together the atoms of reactant and product molecules may be broken while other bonds are formed. Energy is required to break bonds, and since different bonds have different bond dissociation energies, there is often a significant overall energy change in the course of a reaction. In the combustion of methane, for example, all six bonds in the reactant molecules are broken, and six new bonds are formed in the product molecules equation 1.
To analyze such reactions we need to keep track of and evaluate heat changes in a precise and systematic manner. The science that investigates the passage of energy from one system to another, and the transformation of energy from one form to another is called thermodynamics. Classical thermodynamics is a statistical science in which observations are made on macroscopic samples. Heat energy or heat content is designated as enthalpy , symbol H.
In order to consolidate and make use of thermochemical data of this kind, a standard state has been defined. For solutions the standard state is a 1M concentration. The standard states of some typical substances are listed in the following table. Care must be taken to identify the most stable phase or form a given substance will assume under standard conditions. In the case of water this is the liquid phase, and for carbon it is graphite, not diamond.
By definition, the heat of formation of an element in its standard state is 0. A small degree sign to the upper right of the enthalpy symbol e. 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. What mechanism is there for the burnt paper to release its oxygen and return to the original paper state? Even if by some low probability of random states one molecule somewhere did release its oxygen.
That absorbed energy overall, so there is now less energy available locally to kick nearby molecules into doing the same. Systems only settle into their lowest-energy state if they are in thermal equilibrium at zero temperature.
If they are in thermal equilibrium at a nonzero temperature, then they settle into a statistical mixture of all possible energy states that minimizes the Helmholtz free energy. If they aren't in thermal equilibrium, then they can be in any energy state at all, determined by the initial conditions.
Atoms and molecules are coupled with the electromagnetic field. The electrons of the atom and molecule are accelerating and then it will emit electromagnetic radiation , making the system losing energy.
Classical mechanics will predict a total collapse of the atom. Turns out that Quantum Mechanics saves the day, introducing the concept of discrete spectrum and ground state , and the ground state will be a stable state for the system. Quantum mechanics can describe the process of radiation of atoms and molecules, see this for more details of how this works. Now, in general, systems are stable when they minimize the potential energy because there is always a force pushing the system towards the direction that minimizes the potential energy, i.
Atoms and molecules are stable under sufficient small perturbation for other reasons. They are at the ground state, and any perturbation of the system will be reverted into electromagnetic radiation and sent to infinity.
Sign up to join this community. The best answers are voted up and rise to the top. Stack Overflow for Teams — Collaborate and share knowledge with a private group. Create a free Team What is Teams? Learn more. Why are lower energy systems stable?
Ask Question. Asked 4 years, 6 months ago. Active 3 years, 10 months ago. Viewed 4k times. This process of transferring excess energy throughout the metal is attributed to the increase in temperature of the metal. Other ways for the atom to return to stable state is to give off energy in the form of radiation i.
This is not just true of metals. In fact every system wants to minimize its total internal energy and dissipate the extra energy into its surroundings because of the second law of thermodynamics. In the case of an atom, the electromagnetic force is pulling the electrons towards the nucleus, so the further away the electrons from the nucleus, the greater the potential energy of the atom. But there are only certain orbitals or quantum states that electrons can occupy, and they can't occupy the same orbital this is called the Pauli exclusion principle.
So for a given number of electrons, the most stable atomic state is to fill in the orbitals starting from the lowest energy one. If an electron is excited to a higher energy orbital, it will want to quickly go back to a lower available one.
Also, for metals specifically, since they typically have one or two electrons in the outermost shell, if they lose an electron or two they will have lower energy classically an empty or full shell is lower energy than partially full.
Thus they like to form ionic bonds in which they lose electrons. This is called the second law of thermodynamics: any system will move from a more ordered state to a less ordered state. By moving to a lower energy state, the potential energy that was stored in the metal is released to become kinetic energy. Kinetic energy is less ordered than potential energy, and so is favored by thermodynamics.
All systems "want" to be in their lowest energy configuration. Take for example a ball on a hill. A ball on a hill will roll down the hill until it reaches the bottom of the hill. Systems respond to gradients in potential energy be they chemical, gravitational, electrical, or etc. Such gradients produce forces and these forces drive changes in the system. In the example above, there is a gravitational force on the ball due to the change in gravitational potential energy with height gradient in gravitational energy.
If there is no net force acting on a system, then the systems configuration will not change. Such occurrences can be stable, metastable, or unstable.
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