Opening the book…
A reaction is characterized by its equilibrium constant, Keq, a function of the reaction itself, the concentrations of the reactants, and system temperature and pressure. In biological systems we generally ignore pressure (and only occasionally consider temperature), although both can be important for organisms that live on the sea floor, mountain tops, or hydrothermal vents. The equilibrium constant (Keq) for the reaction A + B ⇆ C + D is defined (→) as the concentrations of the products (C + D) divided by the product of the concentrations of the reactants at equilibrium, where nothing macroscopic is happening. At equilibrium the concentrations do not change, that is why K is a constant. For a thermodynamically favorable reaction, one that favors the products, K will be greater, often much greater, than one. The larger Keq, the more product and the less reactant there will be when the system reaches equilibrium. If the equilibrium constant is less than 1 then, at equilibrium the concentration of reactants will be greater than the concentration of products.
\[K=\frac{[C][D]}{[A][B]}\nonumber\]
While the concentration of reactants and products of a reaction at equilibrium remains constant it is not the case that the system is static. If we were to peer into (or imagine) the system at the molecular level we would find that reactants are continuing to form products and products are rearranging to form reactants; the rate of these two reactions will be equal, although slow.[219] Most reactions involve collisions between molecules. The frequency of productive collisions between reactants or products increases as their concentrations increase. At equilibrium for a highly favorable reaction; the high concentration of products (produced by the reaction) x low probability of effective collisions will equal the low concentration of reactants (remaining) x higher probability of effective collisions.
Knowing that a reaction is thermodynamically favorable does not tell us much (or really anything) about whether the reaction occurs to a significant extent under a particular set of conditions. For example, consider a wooden log, which is composed primarily of the carbohydrate polymer cellulose (CH2O)n. In the presence of molecular oxygen (O2) the reaction: \[\mathrm{nO}_2+\log \left(\left(\mathrm{CH}_2 \mathrm{O}\right)_n\right) \leftrightarrows \mathrm{nCO}_2+\mathrm{nH}_2 \mathrm{O}+\text { heat }\nonumber\]is extremely favorable thermodynamically, that is. It has a large negative ΔG and a large equilibrium constant; once the reaction starts it goes completely to CO2 and H2O. Yet logs are stable - they do not spontaneously burst into flames. The question is why not?
The answer lies in the fact that both the equilibrium constant and ΔG (or for the more chemically rigorous, ΔGº) tell us only whether a reaction is thermodynamically favorable; they tell us nothing about whether, or how fast that reaction will proceed under a specific set of conditions. For that we have to turn to the study of reaction rates. This requires us to consider the various factors that affect the reaction. In general a reaction will go faster if there are more reactant molecules. Reactant molecules must collide to initiate a reaction. For example, in the case of the log and oxygen, oxygen molecules (O2) must come in contact with the log. In air (at sea level) O2 molecules amount to ~20% of the total molecules present in the atmosphere. If we increase the atmospheric O2 concentration, the log will burn faster and brighter, because there are more collisions that initiate the reaction.[220] Under normal conditions, however, the log will not start burning spontaneously - added energy is needed. Why? Because the transition between reactants and products requires the breaking of bonds. Breaking a bond requires the addition of energy. Generally, the energy needed is more than the energy available through molecular collisions. The energy required for bonds to break and the reaction to proceed is known as the activation energy. The presence of activation energy explains why chemical systems do not always move spontaneously to equilibrium, why thermodynamically unstable molecules can be "stable". Why nucleic acids and proteins do not quickly breakdown to produce more stable (but rather more boring) molecules such as CO2, H2O, and NH3 from which they are composed.
To explore the idea of activation energy further, let us consider the very simplified model of a log burning in air to produce CO2 and H2O, a reaction that is, in fact, complex. We 
could represent this process on a graph of energy (or more accurately Gibbs Free Energy (G)) vs reaction progress like this (↓). As the reaction proceeds, a great deal of heat is released into the surroundings. This released energy corresponds to the ΔH between reactants and products. The graph also indicates that the products are more stable (lower energy) than the reactants. But, the reaction energy graph does not give us any indication that energy must be added to start the reaction, or why. If we add in this energy the graph would look like this (↓). The activation energy (Ea) is the
energy needed to break the bonds within wood molecules and in O2. These steps, in which pre-existing bonds are broken but new bonds have yet to form is also known as the transition state. In general the amount of activation energy needed to reach the transition state determines the rate of the reaction. If most collisions with surrounding molecules supply this (or more) energy, the reaction will occur rapidly and spontaneously. If, on the other hand and in the case of a log at room temperature, few if any collisions supply enough energy to break the bonds necessary to reach the transition state, the reaction rate will be slow or the reaction will not occur at all (more below). For the wood burning reactions, the energy needed to start the reaction may involve a downed electrical line, a lightning strike, or a burning match.
Once the "forward" reaction starts energy will be released as new bonds are formed. The resulting increase in the temperature (average kinetic energy of molecules) leads to 
increased kinetic energy of colliding reactants; more collisions reach the transition state and produce products. The reaction rate increases and becomes "self-sustaining" - a positive feedback loop (→). As reactants are used up productive collisions between reactants will decrease, while productive collisions between products (the "back reaction") with sufficient activation energy (more energy in fact than the forward reaction) will increase. Unless further fuel (reactants) is added, the reaction system will reach equilibrium. A equilibrium the rates of the forward and back reactions will be equal.
Activation energy and catalysis in biological systems.
As noted above, the reason why (most) thermodynamically favorable reactions do not occur spontaneously when reactants come into contact is that bonds must be broken for the reaction to occur. Breaking bonds, particularly covalent bonds, requires a large amount of energy. In biological systems there are two major sources for this bond-breaking 
energy: light and collisions with other molecules. A molecule can absorb a photon (a particle of light) or energy can be transferred through collisions with other molecules. At room temperature liquid water molecules are moving on average at ~640 meters/second. That is not to say that all molecules are moving at that same speed. If we were to look at the population of molecules, we would find a distribution of speeds known as a Boltzmann (or Maxwell-Boltzmann) distribution (←). As they collide with one another, the molecules exchange energy; one molecule can emerge from a collision with more energy than it entered with. Since reactions occur at temperatures well above absolute zero, there is plenty of energy available in the form of the kinetic energy (movement) of molecules.
Biological systems are constrained in a number of ways. As we will see, the three-dimensional structure of many macromolecules, specifically proteins and nucleic acids, is critical to their normal function, and their 3D structure is often unstable - even small changes, by the standards of a typical chemistry lab, in temperature can lead to what is known as 
denaturation and the loss of function. The take home message is that biological systems use alternative strategies to control the rates of the reactions they depend upon.
Their solution are molecules, typically proteins or small RNAs known as enzymes or ribozymes respectively, that act as catalysts. But what exactly does a catalyst do? Basically, it lowers the "activation" energy of a reaction by interacting with the reactants (→). The result is that, at a particular temperature and environmental conditions, the reaction rate will be increased in the presence of an active catalyst. An important feature of biological catalysts is that their activity can be regulated; their effectiveness as a catalyst for specific reaction can be turned on or off. As we will see, the regulate-ability of biological catalysts is central to maintaining the dynamic, non-equilibrium state of the cell.
Questions to answer:
- How might changes in an organism's internal temperature influence reaction rates?
- A reaction is at equilibrium; we increase the amount of reactant or product. What happens (over time) to the amounts of reactants and products? Drawing some graphs may be helpful.
- What does reducing the activation energy of a reaction do to a system at equilibrium? What does it do to a system far from equilibrium? Drawing some graphs may be helpful.
- How and why does the feedback system of a burning log change over time?
Question to ponder:
- Propose a model for how (at the molecular level) a catalyst might lower a reaction's activation energy?
References
[219] This, of course, assumes that we have a closed system, that is, that neither the products or the reactants can leave the system, and that the volume of the system also remains constant. If the reactants can “leave the scene” of the reaction, then of course the back reaction, Products ⇆Reactants, will be much less likely to occur.
[220] This is one reason why smoking is not a good idea for people who have to use supplementary oxygen to breathe
This page titled 5.3: Reaction rates was last modified on Thu, 08 Jan 2026 23:35:24 GMT and is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Michael W. Klymkowsky and Melanie M. Cooper.
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