# 6.3: Potential, kinetic, free and activation energies (2023)

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##### learning goals

At the end of this section you can do the following:

• define "energy"
• Explain the difference between kinetic energy and potential energy.
• Discuss the concepts of free energy and activation energy.
• Describe endergonic and exergonic reactions.

We define energy as the ability to do work. As you have learned, energy exists in different forms. For example, electrical energy, light energy, and thermal energy are different types of energy. While these are all familiar types of energy that we can see or feel, there is another type of energy that is much less tangible. Scientists associate this energy with something as simple as an object on the ground. To understand how energy enters and leaves biological systems, it is important to understand more about the different types of energy that exist in the physical world.

### types of energy

When an object moves, there is energy. For example, a flying plane generates considerable power. This is because moving objects are capable of making change or doing work. Think of a wrecking ball. Even a slow wrecking ball can do significant damage to other objects. However, a wrecking ball that is not moving cannot do any work. The energy with moving objects isKinetic energy. An accelerating ball, a person walking, the rapid movement of a molecule in air (which generates heat), and electromagnetic radiation such as light all have kinetic energy.

What would happen if we used a crane to lift the same immobile wrecking ball two stories above a car? If the suspended wrecking ball is stationary, can we attribute energy to it? The answer is yes. The suspended wrecking ball has an associated energy that is fundamentally different from the kinetic energy of moving objects. This form of energy results frompotentialfor the wrecking ball to do the work. If we release the ball it will work. Because this type of energy is related to the potential to do work, we call itpotential energy. Objects transfer their energy between kinetic and potential energy as follows: Because the wrecking ball is stationary, it has 0 percent kinetic energy and 100 percent potential energy. Once released, its kinetic energy begins to increase as it accelerates due to gravity. At the same time, as it gets closer to the ground, it loses potential energy. Somewhere in the middle of autumn you have 50% kinetic energy and 50% potential energy. Just before the ball hits the ground, the ball has almost lost its potential energy and is almost at its maximum kinetic energy. Other examples of potential energy include the energy of water trapped by a dam (Figure 6.6) or a person about to jump out of a plane.

Figure 6.6The water behind a dam has potential energy. Moving water, such as a waterfall or rushing river, has kinetic energy. ("dam" Credit: Factory Mod by "Pascal"/Flickr; "Waterfall" Credit: Factory Mod by Frank Gualtieri)

We associate potential energy not only with the location of matter (like a child sitting on a branch), but also with the structure of matter. A spring in the ground has potential energy when it is compressed; as well as a stretched elastic band. The very existence of living cells depends to a large extent on structural potential energy. On a chemical level, the bonds that hold the atoms of molecules together have potential energy. Remember that cellular anabolic pathways require energy to synthesize complex molecules from simpler ones, and catabolic pathways release energy as complex molecules are broken down. That breaking certain chemical bonds can release energy implies that those bonds have potential energy. In fact, every food molecule we eat has potential energy stored in the bonds we eventually use. This is because these bonds can release energy when they break. Scientists name the type of potential energy that exists in chemical bonds and is released when those bonds are broken.chemical energy(Figure 6.7). Chemical energy is responsible for providing living cells with energy from food. The breaking of the molecular bonds within the fuel molecules triggers the release of energy.

Figure 6.7Gasoline molecules contain chemical energy within chemical bonds. This energy is converted to kinetic energy, which allows a car to run on a race track. ("Auto" Credit: Modification of Russell Trow's work)

visit thisPropertyand choose A Simple Pendulum from the menu (under Harmonic Motion) to view the displacement kinetics (K) and potential energy (U) of a moving pendulum.

### free energy

Having learned that chemical reactions release energy when energy storage bonds are broken, the next important question is how we can quantify and express chemical reactions with associated energy. How can we compare the energy released in one reaction with that of another reaction? We use a measure offree energyto quantify these energy transfers. Scientists call this free energy the Gibbs free energy (abbreviated with the letter G) after Josiah Willard Gibbs, the scientist who developed the measurement. According to the second law of thermodynamics, all energy transfers involve the loss of energy in an unusable form, such as energy. B. heat, which leads to entropy. Gibbs free energy refers specifically to the energy involved in a chemical reaction that is available after accounting for entropy. In other words, Gibbs free energy is usable energy, or energy available to do work.

Every chemical reaction involves a change in free energy, called delta G (∆G). We can calculate the free energy change for any system that undergoes such a change, such as B. a chemical reaction. To calculate ∆G, subtract the amount of energy lost to entropy (called ∆S) from the total energy change of the system. The total energy of the system isenthalpyand we denote it as ∆H. The formula to calculate ∆G is as follows, where the symbol T refers to the absolute temperature in Kelvin (degrees Celsius + 273):

$DGRAMS=DH−TDSDGRAMS=DH−TDS$

We express the standard change in free energy of a chemical reaction as the amount of energy per mole of reaction product (in kilojoules or kilocalories, kJ/mol or kcal/mol; 1 kJ = 0.239 kcal) under standard conditions of pH, temperature and pressure. We generally calculate the standard conditions of pH, temperature, and pressure at pH 7.0 in biological systems, 25 degrees Celsius, and 100 kilopascals (1 atm pressure), respectively. Note that cell conditions differ significantly from these standard conditions and therefore the standard ∆G values ​​calculated for biological reactions within the cell will be different.

#### endergonic reactions and exergonic reactions

When energy is released during a chemical reaction, the value resulting from the above equation is a negative number. In other words, reactions that release energy have ∆G < 0. A negative ∆G also means that the reaction products have less free energy than the reactants because they released some free energy during the reaction. Scientists name reactions that have a negative ∆G and therefore release free energy.exergonic reactions. Think:exErgonomic means energy isexitining of the system. We also refer to these reactions as spontaneous reactions because they can occur without adding energy to the system. Understanding which chemical reactions occur spontaneously and release free energy is extremely useful to biologists because these reactions can be used to work within the cell. We have to make an important distinction between the term spontaneous and the idea of ​​a chemical reaction that occurs instantly. Contrary to popular usage of the term, a spontaneous response is not a sudden or quick response. Rusty iron is an example of a spontaneous reaction that occurs slowly, bit by bit, over time.

If a chemical reaction requires an input of energy instead of releasing it, then the ∆G for that reaction is a positive value. In this case, the products have more free energy than the reactants. Therefore, we can think of the products of the reactions as molecules that store energy. We call these chemical reactions.endergonic reactionsand they are not spontaneous. An endergonic reaction does not occur on its own without adding free energy.

Let's go back to the example of the synthesis and breakdown of the food molecule glucose. Remember that building complex molecules like sugar from simpler molecules is an anabolic process and requires energy. Therefore, the chemical reactions involved in anabolic processes are endergonic reactions. Alternatively, the catabolic process of breaking down sugars into simpler molecules releases energy in a series of exergonic reactions. As in the rust example above, the breakdown of sugars involves spontaneous reactions, but these reactions are not instantaneous. Figure 6.8 shows some more examples of endergonic and exergonic reactions. Later sections will provide more information on what else is needed to make even spontaneous reactions more efficient.

### visual connection

#### visual connection

Figure 6.8This figure shows some examples of endergonic processes (those that require energy) and exergonic processes (those that release energy). These include (a) a crumbling compost pile, (b) a chick developing from a fertilized egg, (c) the destruction of sand art, and (d) a ball rolling down a hill. (Credit a: Factory Mod by Natalie Maynor; Credit b: Factory Mod by USDA; Credit c: Factory Mod by "Athlex"/Flickr; Credit d: Factory Mod by Harry Malsch)

Look at each of the processes and decide if it is endergonic or exergonic. In each case, does enthalpy increase or decrease and entropy increase or decrease?

An important concept in the study of metabolism and energy is that of chemical equilibrium. Most chemical reactions are reversible. They can move in both directions, releasing energy to their surroundings in one direction and absorbing energy from the environment in the other direction (Figure 6.9). The same applies to the chemical reactions of cell metabolism, such as the breakdown or building of proteins into individual amino acids. Reactants within a closed system undergo chemical reactions in both directions until they reach an equilibrium state that represents one of the lowest possible free energies and a state of maximum entropy. Energy is required to push the reactants and products out of an equilibrium state. It is necessary to add, remove or change reactants or products. If a cell were a closed system, its chemical reactions would reach equilibrium and it would die because there would not be enough free energy to do the work necessary to sustain life. In a living cell, chemical reactions are constantly moving toward equilibrium but never reaching it. Because a living cell is an open system. Materials go in and out, the cell recycles the products of certain chemical reactions into other reactions, and there is never a chemical equilibrium. In this way, living organisms are involved in an uphill battle, which constantly requires energy, against equilibrium and entropy. This constant supply of energy comes from sunlight which produces nutrients in the process of photosynthesis.

Figure 6.9Both exergonic and endergonic reactions result in changes in the Gibbs free energy (G). Exergonic reactions have a net release of energy and are spontaneous reactions. Endergonic reactions require an input of energy to proceed and are not spontaneous reactions. Both exergonic and endergonic reactions require an initial activation energy for the reaction to occur. Photo credits: Tag, A., Rao, A., Fletcher, S., and Ryan, A. Department of Biology, Texas A&M University.

### activation energy

There is another important concept that we need to keep in mind regarding endergonic and exergonic reactions. Even exergonic reactions require a small energy input before they can continue with their energy-releasing steps. These reactions have a net release of energy but still require some initial energy. Scientists call this small amount of input energy that is necessary for all chemical reactions to take place.activation energy(the energía libre de activación), abbreviated as EA(Figure 6.10).

Why does a reaction releasing negative energy ∆G really need some energy to continue? The reason lies in the steps that take place during a chemical reaction. In chemical reactions, certain chemical bonds are broken and new ones are formed. For example, when a glucose molecule breaks down, the bonds between the carbon atoms in the molecule are broken. Since these are energy storage links, they release energy when broken. However, for them to get to a state where the bonds can be broken, the molecule has to twist a bit. A small power input is required to achieve this contorted state. This distorted state is thattransition state, and is a high-energy unstable state. Because of this, the molecules of the reactants do not stay in their transition state for long, but rather quickly move on to the next steps of the chemical reaction. Free energy diagrams illustrate the energy profiles for a given reaction. Whether the reaction is exergonic or endergonic determine whether the products in the diagram are in a higher or lower energy state than the reactants and products. However, regardless of this measurement, the transition state of the reaction is in a higher energy state than the reactants and therefore EAis always positive.

See an animation of the transition from free energy to the transition state atto beProperty.

Where does the activation energy needed by chemical reagents come from? The necessary source of activation energy to drive the reactions is usually the thermal energy of the environment.Thermal energy(the total binding energy of the reactants or products in a chemical reaction) speeds up the movement of the molecule, increasing the frequency and force with which they collide. It also easily moves atoms and bonds within the molecule, helping them reach their transition state. Because of this, heating a system causes the chemical reactants in that system to react more frequently. Increasing the pressure in a system has the same effect. Once the reactants have absorbed enough thermal energy from their surroundings to reach the transition state, the reaction continues.

The activation energy of a given reaction determines the rate at which it will occur. The higher the activation energy, the slower the chemical reaction. The example of iron oxide illustrates an inherently slow reaction. This reaction is slow over time due to its high EA. Also, the combustion of many fuels, which is highly exergonic, occurs at a negligible rate unless enough heat from a spark exceeds its activation energy. However, once they start to burn, the chemical reactions release enough heat to continue the combustion process and provide the activation energy for the surrounding fuel molecules. As with these extracellular reactions, the activation energy for most cellular reactions is too high for thermal energy to be overcome at efficient rates. In other words, for important cellular reactions to occur at appreciable rates (number of reactions per unit of time), their activation energies must be lowered (Figure 6.10). Scientists refer to this as catalysis. That's a very good thing as far as living cells are concerned. Important macromolecules such as proteins, DNA and RNA store considerable energy and their degradation is exergonic. If cell temperatures alone provided enough thermal energy for these exergonic reactions to overcome their activation barriers, essential cell components would degrade.

### visual connection

#### visual connection

Figure 6.10Activation energy is the energy required for a reaction to occur and is lower when the reaction is catalyzed. The horizontal axis of this diagram describes the chronological sequence of events.

If activation energy was not needed to break down sucrose (table sugar), could you store it in a sugar bowl?

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