Consider the reaction paper + O2→ smoke + ashes + heat + CO2 + H2O Once ignited, the paper burns readily, releasing to the atmosphere both energy as heat and water and carbon dioxide as gases. The reaction is irreversible, since the smoke and ashes never spontaneously retrieve these entities from the heated atmosphere and reconstitute themselves into paper. When the paper burns, its chemical energy is dissipated as heat—not lost from the universe, since energy can never be created or destroyed, but irretrievably dispersed in the chaotic random thermal motions of molecules. At the same time, the atoms and molecules of the paper become dispersed and disordered. In the language of thermodynamics, there has been a loss of free energy; that is, of energy that can be harnessed to do work or drive chemical reactions. This loss reflects a reduction of orderliness in the way the energy and molecules were stored in the paper. We shall discuss free energy in more detail shortly, but the general principle is clear enough intuitively: chemical reactions proceed spontaneously only in the direction that leads to a loss of free energy. In other words, the spontaneous direction for any reaction is the direction that goes “downhill,” where a “downhill” reaction is one that is energetically favorable. Although the most energetically favorable form of carbon under ordinary conditions is CO2, and that of hydrogen is H2O, a living organism does not disappear in a puff of smoke, and the paper book in your hands does not burst into flames. This is because the molecules both in the living organism and in the book are in a relatively stable state, and they cannot be changed to a state of lower energy without an input of energy: in other words, a molecule requires activation energy—a kick over an energy barrier—before it can undergo a chemical reaction that leaves it in a more stable state (Figure 2–21). In the case of a burning book, the activation energy can be provided by the heat of a lighted match. For the molecules in the watery solution inside a cell, the kick is delivered by an unusually energetic random collision with surrounding molecules—collisions that become more violent as the temperature is raised. The chemistry in a living cell is tightly controlled, because the kick over energy barriers is greatly aided by a specialized class of proteins—the enzymes. Each enzyme binds tightly to one or more molecules, called substrates, and holds them in a way that greatly reduces the activation energy of a particular chemical reaction that the bound substrates can undergo. A substance that can lower the activation energy of a reaction is termed a catalyst; catalysts increase the rate of chemical reactions because they allow a much larger proportion of the random collisions with surrounding molecules to kick the substrates over the energy barrier, as illustrated in Figure 2–22. Enzymes are among the most effective catalysts
Figure 2–21 The important principle of activation energy. (a) Compound Y (a reactant) is in a relatively stable state, and energy is required to convert it to compound X (a product), even though X is at a lower overall energy level than Y. This conversion will not take place, therefore, unless compound Y can acquire enough activation energy (energy a minus energy b) from its surroundings to undergo the reaction that converts it into compound X. This energy may be provided by means of an unusually energetic collision with other molecules. for the reverse reaction,  X → Y, the activation energy will be d more rarely. activation energies are always positive; note, however, that the total energy change for the energetically known: some are capable of speeding up reactions by factors of 1014 or more. Enzymes thereby allow reactions that would not otherwise occur to proceed rapidly at normal temperatures.
An enzyme cannot change the equilibrium point for a reaction. The reason is simple: when an enzyme (or any catalyst) lowers the activation energy for the reaction Y → X, of necessity it also lowers the activation energy for the reaction X → Y by exactly the same amount (see Figure 2–21). The forward and backward reactions will therefore be accelerated by the same factor by an enzyme, and the equilibrium point for the reaction will be unchanged (Figure 2–23). Thus no matter how much an enzyme speeds up a reaction, it cannot change its direction.