Let's start to think a little about the rate of a reaction. Even exergonic (spontaneous) reactions often require a small increase in free energy before they can begin converting reactants to products. Hydrogen and Oxygen gas do not react when mixed at room temperature, they need a spark, which energizes some of the molecules, giving them the kick they need to react. After the initial spark, the heat released by this exergonic and exothermic reaction continues to provide the kick needed to get H2 and O2 to react to form water. This small positive change in free energy required to initiate the reaction, first provided by the spark, and then by the ongoing reaction, is called the activation energy (or free energy of activation) and is sometimes abbreviated EA. Not all exergonic reactions need to be heated up to proceed- the reaction of sodium metal and chlorine gas proceeds immediately and enthusiastically on mixing. However, almost any biologically relevant reaction, while exergonic (or it would not proceed at all, right?), does not proceed in an uncontrolled fashion. This may initially seem odd, as we don't have anything like a spark plug in our cells. As we'll see soon, we employ protein-based catalysts, called enzymes, to regulate reaction rate. But here, let's discuss activation energy, its relationship to the transition state, and what catalysts do.
The oxidation of gasoline is highly exergonic. Despite this, why do cars not spontaneously explode in parking lots?
Why do many chemical reactions - even those with a very large negative ∆G - first require some energy increase to proceed? The reason lies in the steps that take place during a chemical reaction. Chemical reactions, almost by definition, require that some chemical bonds be broken and/or formed. For example, when a glucose molecule is broken down, the glycosidic bonds are broken, bonds within water are broken and new bonds are made between the "disassembled" water and the atoms that were involved in the glycosidic bond. While the overall reaction (the combination of energy cost of breaking bonds, energy gained by making new bonds and the change of entropy between reactants and products) may have a negative ∆G (which means it can proceed... eventually) the breaking of the bonds, or the formation of very unusual and unstable intermediate structures, requires some level of energy "investment". This structural energy invested in highly unstable intermediates will be released again when the intermediate breaks down (either to re-form the reactants, or instead form the products). The configuration of this very unstable intermediate is termed the transition state. The height of the activation energy "barrier" has a relationship to the rate of a reaction. The higher/larger the barrier, the slower the reaction. Why is this? Looking at the diagram below (a typical diagram used to illustrate the activation energy concept), we would guess that, although energetically favorable, this reaction would never proceed the reactants do not have the potential energy required. The answer is that energy is present not just in the structure of the molecules, but also in their kinetic energy, which takes the form of both vibration (within the molecule) and speed (with which the molecule is moving). These energies (especially in a gas, but also in liquid) are fairly narrowly distributed within the population of molecules. But sometimes we might imagine that a molecule might get lucky- perhaps it is struck by several other molecules in a short space of time, and their energy is conferred to it. That one lucky molecule might reach the energy required to react. This might be the scenario for a exergonic, but slow, reaction. Or a chemist could heat up the molecules, the potential energy of all the reactants (and products) molecules would rise, and the majority of molecules could reach the energy required for the transition state, and the reaction could proceed quickly.
In the diagram below, the "G" (Potential energy) of the reactants is described as a single value on the Y axis. Is this true? If you think it isn't true, how would you "correct" the diagram?
We said the heating the reaction would increase the G of the reactants and products. Would it also increase the value of the transition state?
Can you propose a physical analog (or model) that can help explain why the activation energy barrier is related to the rate of the reaction, whereas the free energy difference between substrate and product is not.
Activation energy is the energy required for a reaction to proceed, and it is lower if the reaction is catalyzed. This Y axis of this diagram tells us the potential energy of the reactants and products (which is based on their molecular structure, concentration, and, in this case, since we're discussing sparks, we'll consider temperature too- usually we regard that as a constant). The potential energy of the transition state is based only on its structure. The X axis here is virtually meaningless. It certainly does not have units of "time"- the entire set of molecules is not moving from left to right simultaneously!
As noted, the activation energy of a particular reaction determines the rate at which it will proceed. The higher the activation energy, the slower the chemical reaction will be. In fact, in a solid or liquid, the distribution of energies between molecules is pretty uniform; if the activation energy is significant above the energy of the reactants the reaction will not occur at a biologically significant rate. The example of iron rusting illustrates an inherently slow reaction. The conversion of diamond into graphite is another spontaneous reaction that take a LONG time. These reactions occur slowly over time because of high activation energy barriers. The burning (oxidation) of many fossil fuels, which is an exergonic and exothermic process, will take place at a negligible rate unless their activation energy is overcome by sufficient heat from a spark or match. Once these fuels begin to burn, however, the chemical reactions release enough heat to help overcome the activation energy barrier for the combustion of the rest of the fuel. Like these reactions outside of cells, the activation energy for most cellular reactions is too high for heat energy to overcome at efficient rates. By the way, this is a very good thing as far as living cells are concerned. Important macromolecules, such as proteins, DNA, and RNA, store considerable energy, and their breakdown is exergonic. If cellular temperatures alone provided enough heat energy for these exergonic reactions to overcome their activation barriers, the essential components of a cell would disintegrate. Therefore, in order for important cellular reactions to occur at appreciable rates (number of reactions per unit time), their activation energies, somehow, must be lowered via some other approach- not by burning!
Catalysts lower activation energy by providing an alternate reaction pathway- a route from reactant to product that does not require a high-energy intermediate. Many metals are excellent catalysts as they are intrinsically flexible in their oxidation state, and so are adept at taking and then releasing electrons (in other words, acting as an intermediate that helps rearrange bonds between other molecules). It is important to bear in mind that catalysts take, and then release (or else give, then receive), electrons. Thus means that during the progress of the reaction, the catalyst is regenerated. It participates in the reaction, but comes out the reaction in its original state, ready to participate again. Life regulates reaction rate using proteinaceous catalysts called enzymes- these often carry metals that perform the catalysis, while the protein's shape restricts the types of reactants and products that can access that metal, thus providing specificity to the reaction.
If no activation energy were required to break down sucrose (table sugar), would you be able to store it in a sugar bowl?
Lowering the activation energy:
- makes the reaction happen faster
- lowers the energy level of the transition state
- is accomplished by adding a catalyst to the reaction
- always causes more product to be produced
- only reduces the transition state energy level in one direction, from reactants to products
- a, b and c
- b and c
- all of the above are true
Which of the following comparisons or contrasts between endergonic and exergonic reactions is false?
- Endergonic reactions have a positive ∆G and exergonic reactions have a negative ∆G
- Endergonic reactions consume energy and exergonic reactions release energy
- Both endergonic and exergonic reactions require a small amount of energy to overcome an activation barrier
- Endergonic reactions take place slowly and exergonic reactions take place quickly
Appendix: Energy Units
In the International System of Units (SI), the unit of work or energy is the Joule (J). For very small amounts of energy, the erg (erg) is sometimes used. An erg is one ten millionth of a Joule:
1 Joule=10,000,000 ergs
Heat energy is often measured in calories. One calorie (cal) is defined as the heat required to raise the temperature of 1 gram of water from 14.5 to 15.5 ºC. 1 calorie = 4.189 Joules
Extremely annoying is the definition of Calorie (Cal, note the subtle difference in abbreviation), often used in nutrition, which is the energy required to raise the temperature of a kilogram of water by 1 degree.
In Biology, we'll talk about energy a lot, but we really don't think much about power. But here are some units for your enjoyment.
Power is the rate at which energy is used. The unit of power is the Watt (W), named after James Watt, who perfected the steam engine:
1 Watt=1 Joule/second
Electrical energy is generally expressed in kilowatt-hours (kWh):
1 kilowatt−hour=3,600,000 Joules