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Energy in chemical reactions
Chemical reactions involve a redistribution of energy within the reacting chemicals and with their environment. So, like it or not, we need to develop some models that can help us to describe where energy is in a system (perhaps how it is "stored"/distrbuted) and how it can be moved around in a reaction. The models we develop will not be overly detailed in the sense that they would satisfy a hard-core chemist or physicist with their level of technical detail, but we expect that they should still be technically correct and not form incorrect mental models that will make it difficult to understand the "refinements" later.
In this respect, one of the key concepts to understand is that we are going to think about energy being transferred between parts of a system. We'll try not to think about it as being transformed. The distinction between "transfer" and "transform" is important. The latter gives the impression that energy is a property that exists in different forms, that it gets reshaped somehow. One problem with the "transform" language is that it is difficult to reconcile with the idea that energy is being conserved (according to the first law of thermodynamics) if it is constantly changing form. How can the entity be conserved if it is no longer the same thing? Moreover, the second law of thermodynamics tells us that no transformation conserves all energy in a system. If energy is getting "transformed," how can it be conserved?
So, instead, we are going to approach this issue by transferring and storing energy between different parts of a system and thus think about energy as a property that can get redistributed. That'll hopefully make the accounting of energy easier.
If we are going to think about transferring energy from one part of a system to another, we also need to be careful about NOT treating energy like a substance that moves like a fluid or "thing." Rather, we need to appreciate energy simply as a property of a system that can be measured and reorganized but that is neither a "thing" nor something that is at one time in one form then later in another.
Since we will often be dealing with transformations of biomolecules, we can start by thinking about where energy can be found/stored in these systems. We'll start with a couple of ideas and add more to them later.
Let us propose that one place that energy can be stored is in the motion of matter. For brevity, we'll give the energy stored in motion a name: kinetic energy. Molecules in biology are in constant motion and therefore have a certain amount of kinetic energy (energy stored in motion) associated with them.
Let us also propose that there is a certain amount of energy stored in the biomolecules themselves and that the amount of energy stored in those molecules is associated with the types and numbers of atoms in the molecules and their organization (the number and types of bonds between them). The discussion of exactly where the energy is stored in the molecules is beyond the scope of this class, but we can approximate it by suggesting that a good proxy is in the bonds. Different types of bonds may be associated with storing different amounts of energy. In some contexts, this type of energy storage could be labeled potential energy or chemical energy. With this view, one of the things that happens during the making and breaking of bonds in a chemical reaction is that the energy is transferred about the system into different types of bonds. In the context of an Energy Story (a topic in another module), one could theoretically count the amount of energy stored in the bonds and motion of the reactants and the energy stored in the bonds and energy of the products.
In some cases, you might find that when you add up the energy stored in the products and the energy stored in the reactants that these sums are not equal. If the energy in the reactants is greater than that in the products, where did this energy go? It had to get transferred to something else. Some will certainly have moved into other parts of the system, stored in the motion of other molecules (warming the environment) or perhaps in the energy associated with photons of light. One good, real-life example is the chemical reaction between wood and oxygen (reactants) and it's conversion to carbon dioxide and water (products). At the beginning, the energy in the system is largely in the molecular bonds of oxygen and the wood (reactants). There is still energy left in the carbon dioxide and water (products) but less than at the beginning. We all appreciate that some of that energy was transferred to the energy in light and heat. This reaction where energy is transferred to the environment is termed exothermic. By contrast, in some reactions, energy will transfer in from the environment. These reactions are endothermic.
The transfer of energy in or out of the reaction from the environment is NOT the only thing that determines whether a reaction will be spontaneous or not. We'll discuss that soon. For the moment, it is important to get comfortable with the idea that energy can be transferred among different components of a system during a reaction and that you should be able to envision tracking it.
Energy and Chemical Reactions# - Biology
Scientists use the term bioenergetics to describe the concept of energy flow (Figure 1) through living systems, such as cells. Cellular processes such as the building and breaking down of complex molecules occur through stepwise chemical reactions. Some of these chemical reactions are spontaneous and release energy, whereas others require energy to proceed.
Figure 1. Ultimately, most life forms get their energy from the sun. Plants use photosynthesis to capture sunlight, and herbivores eat the plants to obtain energy. Carnivores eat the herbivores, and eventual decomposition of plant and animal material contributes to the nutrient pool.
Just as living things must continually consume food to replenish their energy supplies, cells must continually produce more energy to replenish that used by the many energy-requiring chemical reactions that constantly take place. Together, all of the chemical reactions that take place inside cells, including those that consume or generate energy, are referred to as the cell’s metabolism.
- Identify different types of metabolic pathways
- Distinguish between an open and a closed system
- State the first law of thermodynamics
- State the second law of thermodynamics
- Explain the difference between kinetic and potential energy
- Describe endergonic and exergonic reactions
- Discuss how enzymes function as molecular catalysts
This team of ants is breaking down a dead tree. A classic example of teamwork. And all that work takes energy. In fact, each chemical reaction - the chemical reactions that allow the cells in those ants to do the work - needs energy to get started. And all that energy comes from the food the ants eat. Whatever eats the ants gets their energy from the ants. Energy passes through an ecosystem in one direction only.
Chemical reactions always involve energy. Energy is a property of matter that is defined as the ability to do work. When methane burns, for example, it releases energy in the form of heat and light. Other chemical reactions absorb energy rather than release it.
A chemical reaction that releases energy (as heat) is called an exothermic reaction. This type of reaction can be represented by a general chemical equation:
Reactants &rarr Products + Heat
In addition to methane burning, another example of an exothermic reaction is chlorine combining with sodium to form table salt. This reaction also releases energy.
A chemical reaction that absorbs energy is called an endothermic reaction. This type of reaction can also be represented by a general chemical equation:
Reactants + Heat &rarr Products
Did you ever use a chemical cold pack? The pack cools down because of an endothermic reaction. When a tube inside the pack is broken, it releases a chemical that reacts with water inside the pack. This reaction absorbs heat energy and quickly cools down the pack.
All chemical reactions need energy to get started. Even reactions that release energy need a boost of energy in order to begin. The energy needed to start a chemical reaction is called activation energy. Activation energy is like the push a child needs to start going down a playground slide. The push gives the child enough energy to start moving, but once she starts, she keeps moving without being pushed again. Activation energy is illustrated in Figure below.
Activation Energy. Activation energy provides the &ldquopush&rdquo needed to start a chemical reaction. Is the chemical reaction in this figure an exothermic or endothermic reaction?
Why do all chemical reactions need energy to get started? In order for reactions to begin, reactant molecules must bump into each other, so they must be moving, and movement requires energy. When reactant molecules bump together, they may repel each other because of intermolecular forces pushing them apart. Overcoming these forces so the molecules can come together and react also takes energy.
Potential and Kinetic Energy
When an object is in motion, there is energy associated with that object. Think of a wrecking ball. Even a slow-moving wrecking ball can do a great deal of damage to other objects. Energy associated with objects in motion is called kinetic energy (Figure 5). A speeding bullet, a walking person, and the rapid movement of molecules in the air (which produces heat) all have kinetic energy.
Now what if that same motionless wrecking ball is lifted two stories above ground with a crane? If the suspended wrecking ball is unmoving, is there energy associated with it? The answer is yes. The energy that was required to lift the wrecking ball did not disappear, but is now stored in the wrecking ball by virtue of its position and the force of gravity acting on it. This type of energy is called potential energy (Figure 5). If the ball were to fall, the potential energy would be transformed into kinetic energy until all of the potential energy was exhausted when the ball rested on the ground. Wrecking balls also swing like a pendulum through the swing, there is a constant change of potential energy (highest at the top of the swing) to kinetic energy (highest at the bottom of the swing). Other examples of potential energy include the energy of water held behind a dam or a person about to skydive out of an airplane.
Figure 5 Still water has potential energy moving water, such as in a waterfall or a rapidly flowing river, has kinetic energy. (credit “dam”: modification of work by “Pascal”/Flickr credit “waterfall”: modification of work by Frank Gualtieri)
Potential energy is not only associated with the location of matter, but also with the structure of matter. A spring on the ground has potential energy if it is compressed so does a rubber band that is pulled taut. On a molecular level, the bonds that hold the atoms of molecules together exist in a particular structure that has potential energy. Cellular pathways require energy to synthesize complex molecules from simpler ones and other pathways release energy when these complex molecules are broken down. The fact that energy can be released by the breakdown of certain chemical bonds implies that those bonds have potential energy. In fact, there is potential energy stored within the bonds of all the food molecules we eat, which is eventually harnessed for use. This is because these bonds can release energy when broken. The type of potential energy that exists within chemical bonds, and is released when those bonds are broken, is called chemical energy. Chemical energy is responsible for providing living cells with energy from food. The release of energy occurs when the molecular bonds within food molecules are broken.
Specific heat capacity (c) - the quantity of thermal energy required to raise the temperature of 1 g of a substance by 1 °C SI units – J/(g∙°C).
Substances with a high specific heat capacity take longer to heat or to cool.
Specific Heat Capacities of Some Common Substances
The ΔG of a Reaction Depends on Changes in Enthalpy (Bond Energy) and Entropy
At any constant temperature and pressure, two factors determine the ΔG of a reaction and thus whether the reaction will tend to occur: the change in bond energy between reactants and products and the change in the randomness of the system. Gibbs showed that free energy can be defined as
Entropy S is a measure of the degree of randomness or disorder of a system. Entropy increases as a system becomes more disordered and decreases as it becomes more structured. Consider, for example, the diffusion of solutes from one solution into another one in which their concentration is lower. This important biological reaction is driven only by an increase in entropy in such a process ΔH is near zero. To see this, suppose that a 0.1 M solution of glucose is separated from a large volume of water by a membrane through which glucose can diffuse. Diffusion of glucose molecules across the membrane will give them more room in which to move, with the result that the randomness, or entropy, of the system is increased. Maximum entropy is achieved when all molecules can diffuse freely over the largest possible volume — that is, when the concentration of glucose molecules is the same on both sides of the membrane. If the degree of hydration of glucose does not change significantly on dilution, ΔH will be approximately zero the negative free energy of the reaction in which glucose molecules are liberated to diffuse over a larger volume will be due solely to the positive value of ΔS in Equation 2-7.
As mentioned previously, the formation of hydrophobic bonds is driven primarily by a change in entropy. That is, if a long hydrophobic molecule, such as heptane or tristearin, is dissolved in water, the water molecules are forced to form a cage around it, restricting their free motion. This imposes a high degree of order on their arrangement and lowers the entropy of the system (ΔS <𠁐). Because the entropy change is negative, hydrophobic molecules do not dissolve well in aqueous solutions and tend to stay associated with one another.
We can summarize the relationships between free energy, enthalpy, and entropy as follows:
Many biological reactions lead to an increase in order, and thus a decrease in entropy (ΔS <𠁐). An obvious example is the reaction that links amino acids together to form a protein. A solution of protein molecules has a lower entropy than does a solution of the same amino acids unlinked, because the free movement of any amino acid in a protein is restricted when it is bound in a long chain. For the linking reaction to proceed, a compensatory decrease in free energy must occur elsewhere in the system, as is discussed in Chapter 4.
After learning that chemical reactions release energy when energy-storing bonds are broken, an important next question is the following: How is the energy associated with these chemical reactions quantified and expressed? How can the energy released from one reaction be compared to that of another reaction? A measurement of free energy is used to quantify these energy transfers. Recall that according to the second law of thermodynamics, all energy transfers involve the loss of some amount of energy in an unusable form such as heat. Free energy specifically refers to the energy associated with a chemical reaction that is available after the losses are accounted for. In other words, free energy is usable energy, or energy that is available to do work.
If energy is released during a chemical reaction, then the change in free energy, signified as ∆G (delta G) will be a negative number. A negative change in free energy also means that the products of the reaction have less free energy than the reactants, because they release some free energy during the reaction. Reactions that have a negative change in free energy and consequently release free energy are called exergonic reactions. Think: exergonic means energy is exiting the system. These reactions are also referred to as spontaneous reactions, and their products have less stored energy than the reactants. An important distinction must be drawn between the term spontaneous and the idea of a chemical reaction occurring immediately. Contrary to the everyday use of the term, a spontaneous reaction is not one that suddenly or quickly occurs. The rusting of iron is an example of a spontaneous reaction that occurs slowly, little by little, over time.
If a chemical reaction absorbs energy rather than releases energy on balance, then the ∆G for that reaction will be a positive value. In this case, the products have more free energy than the reactants. Thus, the products of these reactions can be thought of as energy-storing molecules. These chemical reactions are called endergonic reactions and they are non-spontaneous. An endergonic reaction will not take place on its own without the addition of free energy.
Figure 4.6 Shown are some examples of endergonic processes (ones that require energy) and exergonic processes (ones that release energy). (credit a: modification of work by Natalie Maynor credit b: modification of work by USDA credit c: modification of work by Cory Zanker credit d: modification of work by Harry Malsch)
Look at each of the processes shown and decide if it is endergonic or exergonic.
There is another important concept that must be considered regarding endergonic and exergonic reactions. Exergonic reactions require a small amount of energy input to get going, before they can proceed with their energy-releasing steps. These reactions have a net release of energy, but still require some energy input in the beginning. This small amount of energy input necessary for all chemical reactions to occur is called the activation energy.
Independent Practice - Modeling The Energy Of Photosynthesis
Students will use the Lecture Notes from the previous section as a resource to develop a model that illustrates the light and dark (Calvin Cycle) reactions of photosynthesis. Students do not need to get bogged down with the details of biochemical steps of each process. The focus should be placed on the inputs (starting materials) and the outputs (ending materials) for each of the two steps. Students are also encouraged to follow the transformation of energy as it is absorbed as sunlight (light energy) and is converted into glucose (stored chemical energy) for the plant to use during cellular respiration.
Common Student Misconceptions :
- Plants can only undergo photosynthesis and will not experience cellular respiration. **Plants experience cellular respiration in the mitochondria which will convert the stored glucose into energy to grow and sustain life for the plant.
- The dark reaction cannot occur in the light. **The dark reaction (Calvin Cycle) is able occur in the light, but is only called the dark reaction because it is light-independent, meaning the this chemical reaction does not NEED the sunlight to occur.
- The chloroplast is the only organelle to undergo chemical reactions. **All of the plant's organelles are undergoing chemical reactions. The chloroplast contains chlorophyll that absorb the sun's energy that powers the chemical reaction of photosynthesis.
- Photosynthesis will occur in a plant no matter what! **The chemical reaction of photosynthesis need the sunlight and the chlorophyll pigment to start the chemical reaction in the thylakoid membrane. If there is not sunlight (or artificial UV light) then the process of photosynthesis cannot occur.
- The Calvin Cycle can occur by itself without the light reaction. **The Calvin Cycle needs the products of the light reaction to occur,so the light reaction must occur for the Calvin Cycle to proceed. The light reaction provides the materials to allow the dark reaction (Calvin Cycle) to occur.
Sample of Student Work: Illustrated Models and Narrations
Sample of Student Work: Exemplary Photosynthesis Diagram - This artifact demonstrates the student's attention to detail while trying to master the intricate chemical processes associated with photosynthesis. The student's effort will support her learning as she examines the diagram to follow each phase of the process as sunlight is converted to stored chemical energy.
Sample of Student Work: Needs Improvement Photosynthesis Diagram- This artifact displays a student's work that appears to be rushed with very little effort. The greatest concern is that the illustrated model is difficult to read so the student would have a tough time going back to study the model while trying to prepare for an assessment. Students do not have to be talented artists to be successful in this assignment, but they do need to have attention to detail and attempt to make their illustrated model as neat as possible in order to support their learning of the content.
Sample of Student Work: Narration for the Process of Photosynthesis:This student's summary narration of photosynthesis demonstrates a basic understanding of the chemical processes of this complicated chemical reaction. As the unit of study progresses, the additional lessons will strengthen this student's level of comprehension.
28 Potential, Kinetic, Free, and Activation Energy
By the end of this section, you will be able to do the following:
- Define “energy”
- Explain the difference between kinetic 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’ve learned, energy exists in different forms. For example, electrical energy, light energy, and heat energy are all different energy types. While these are all familiar energy types that one can see or feel, there is another energy type that is much less tangible. Scientists associate this energy with something as simple as an object above the ground. In order to appreciate the way energy flows into and out of biological systems, it is important to understand more about the different energy types that exist in the physical world.
When an object is in motion, there is energy. For example, an airplane in flight produces considerable energy. This is because moving objects are capable of enacting a change, or doing work. Think of a wrecking ball. Even a slow-moving wrecking ball can do considerable damage to other objects. However, a wrecking ball that is not in motion is incapable of performing work. Energy with objects in motion is kinetic energy . A speeding bullet, a walking person, rapid molecule movement in the air (which produces heat), and electromagnetic radiation like light all have kinetic energy.
What if we lift that same motionless wrecking ball two stories above a car with a crane? If the suspended wrecking ball is unmoving, can we associate energy with it? The answer is yes. The suspended wrecking ball has associated energy that is fundamentally different from the kinetic energy of objects in motion. This energy form results from the potential for the wrecking ball to do work. If we release the ball it would do work. Because this energy type refers to the potential to do work, we call it potential energy . Objects transfer their energy between kinetic and potential in the following way: As the wrecking ball hangs motionless, it has 0 kinetic and 100 percent potential energy. Once it releases, its kinetic energy begins to increase because it builds speed due to gravity. Simultaneously, as it nears the ground, it loses potential energy. Somewhere mid-fall it has 50 percent kinetic and 50 percent potential energy. Just before it hits the ground, the ball has nearly lost its potential energy and has near-maximal kinetic energy. Other examples of potential energy include water’s energy held behind a dam ((Figure)), or a person about to skydive from an airplane.
We associate potential energy only with the matter’s location (such as a child sitting on a tree branch), but also with the matter’s structure. A spring on the ground has potential energy if it is compressed so does a tautly pulled rubber band. The very existence of living cells relies heavily on structural potential energy. On a chemical level, the bonds that hold the molecules’ atoms together have potential energy. Remember that anabolic cellular pathways require energy to synthesize complex molecules from simpler ones, and catabolic pathways release energy when complex molecules break down. That certain chemical bonds’ breakdown can release energy implies that those bonds have potential energy. In fact, there is potential energy stored within the bonds of all the food molecules we eat, which we eventually harness for use. This is because these bonds can release energy when broken. Scientists call the potential energy type that exists within chemical bonds that releases when those bonds break chemical energy ((Figure)). Chemical energy is responsible for providing living cells with energy from food. Breaking the molecular bonds within fuel molecules brings about the energy’s release.
Visit this site and select “A simple pendulum” on the menu (under “Harmonic Motion”) to see the shifting kinetic (K) and potential energy (U) of a pendulum in motion.
After learning that chemical reactions release energy when energy-storing bonds break, an important next question is how do we quantify and express the chemical reactions with the associated energy? How can we compare the energy that releases from one reaction to that of another reaction? We use a measurement of free energy to quantitate these energy transfers. Scientists call this free energy Gibbs free energy (abbreviated with the letter G) after Josiah Willard Gibbs, the scientist who developed the measurement. Recall that according to the second law of thermodynamics, all energy transfers involve losing some energy in an unusable form such as heat, resulting in entropy. Gibbs free energy specifically refers to the energy that takes place with a chemical reaction that is available after we account for entropy. In other words, Gibbs free energy is usable energy, or energy that is available to do work.
Every chemical reaction involves a change in free energy, called delta G (∆G). We can calculate the change in free energy for any system that undergoes such a change, such as a chemical reaction. To calculate ∆G, subtract the amount of energy lost to entropy (denoted as ∆S) from the system’s total energy change. Scientists call this total energy change in the system enthalpy and we denote it as ∆H. The formula for calculating ∆G is as follows, where the symbol T refers to absolute temperature in Kelvin (degrees Celsius + 273):
We express a chemical reaction’s standard free energy change as an amount of energy per mole of the reaction product (either in kilojoules or kilocalories, kJ/mol or kcal/mol 1 kJ = 0.239 kcal) under standard pH, temperature, and pressure conditions. We generally calculate standard pH, temperature, and pressure conditions at pH 7.0 in biological systems, 25 degrees Celsius, and 100 kilopascals (1 atm pressure), respectively. Note that cellular conditions vary considerably from these standard conditions, and so standard calculated ∆G values for biological reactions will be different inside the cell.
Endergonic Reactions and Exergonic Reactions
If energy releases during a chemical reaction, then the resulting value from the above equation will be a negative number. In other words, reactions that release energy have a ∆G < 0. A negative ∆G also means that the reaction’s products have less free energy than the reactants, because they gave off some free energy during the reaction. Scientists call reactions that have a negative ∆G and consequently release free energy exergonic reactions . Think: exergonic means energy is exiting the system. We also refer to these reactions as spontaneous reactions, because they can occur without adding energy into the system. Understanding which chemical reactions are spontaneous and release free energy is extremely useful for biologists, because these reactions can be harnessed to perform work inside the cell. We must draw an important distinction between the term spontaneous and the idea of a chemical reaction that occurs immediately. Contrary to the everyday use of the term, a spontaneous reaction is not one that suddenly or quickly occurs. Rusting iron is an example of a spontaneous reaction that occurs slowly, little by little, over time.
If a chemical reaction requires an energy input rather than releasing energy, then the ∆G for that reaction will be a positive value. In this case, the products have more free energy than the reactants. Thus, we can think of the reactions’ products as energy-storing molecules. We call these chemical reactions endergonic reactions , and they are non-spontaneous. An endergonic reaction will not take place on its own without adding free energy.
Let’s revisit the example of the synthesis and breakdown of the food molecule, glucose. Remember that building complex molecules, such as sugars, from simpler ones is an anabolic process and requires energy. Therefore, the chemical reactions involved in anabolic processes are endergonic reactions. Alternatively the catabolic process of breaking sugar down into simpler molecules releases energy in a series of exergonic reactions. Like the rust example above, the sugar breakdown involves spontaneous reactions, but these reactions do not occur instantaneously. (Figure) shows some other examples of endergonic and exergonic reactions. Later sections will provide more information about what else is required to make even spontaneous reactions happen more efficiently.
Look at each of the processes, and decide if it is endergonic or exergonic. In each case, does enthalpy increase or decrease, and does entropy increase or decrease?
An important concept in studying metabolism and energy is that of chemical equilibrium. Most chemical reactions are reversible. They can proceed in both directions, releasing energy into their environment in one direction, and absorbing it from the environment in the other direction ((Figure)). The same is true for the chemical reactions involved in cell metabolism, such as the breaking down and building up of proteins into and from individual amino acids, respectively. Reactants within a closed system will undergo chemical reactions in both directions until they reach a state of equilibrium, which is one of the lowest possible free energy and a state of maximal entropy. To push the reactants and products away from a state of equilibrium requires energy. Either reactants or products must be added, removed, or changed. If a cell were a closed system, its chemical reactions would reach equilibrium, and it would die because there would be insufficient free energy left to perform the necessary work to maintain life. In a living cell, chemical reactions are constantly moving towards equilibrium, but never reach it. This is because a living cell is an open system. Materials pass in and out, the cell recycles the products of certain chemical reactions into other reactions, and there is never chemical equilibrium. In this way, living organisms are in a constant energy-requiring, uphill battle against equilibrium and entropy. This constant energy supply ultimately comes from sunlight, which produces nutrients in the photosynthesis process.
There is another important concept that we must consider regarding endergonic and exergonic reactions. Even exergonic reactions require a small amount of energy input before they can proceed 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 energy input necessary for all chemical reactions to occur the activation energy (or free energy of activation) abbreviated as EA ((Figure)).
Why would an energy-releasing, negative ∆G reaction actually require some energy to proceed? The reason lies in the steps that take place during a chemical reaction. During chemical reactions, certain chemical bonds break and new ones form. For example, when a glucose molecule breaks down, bonds between the molecule’s carbon atoms break. Since these are energy-storing bonds, they release energy when broken. However, to get them into a state that allows the bonds to break, the molecule must be somewhat contorted. A small energy input is required to achieve this contorted state. This contorted state is the transition state , and it is a high-energy, unstable state. For this reason, reactant molecules do not last long in their transition state, but very quickly proceed to the chemical reaction’s next steps. Free energy diagrams illustrate the energy profiles for a given reaction. Whether the reaction is exergonic or endergonic determines whether the products in the diagram will exist at a lower or higher energy state than both the reactants and the products. However, regardless of this measure, the transition state of the reaction exists at a higher energy state than the reactants, and thus, EA is always positive.
Watch an animation of the move from free energy to transition state at this site.
From where does the activation energy that chemical reactants require come? The activation energy’s required source to push reactions forward is typically heat energy from the surroundings. Heat energy (the total bond energy of reactants or products in a chemical reaction) speeds up the molecule’s motion, increasing the frequency and force with which they collide. It also moves atoms and bonds within the molecule slightly, helping them reach their transition state. For this reason, heating a system will cause chemical reactants within that system to react more frequently. Increasing the pressure on a system has the same effect. Once reactants have absorbed enough heat energy from their surroundings to reach the transition state, the reaction will proceed.
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. The example of iron rusting illustrates an inherently slow reaction. This reaction occurs slowly over time because of its high EA. Additionally, burning many fuels, which is strongly exergonic, will take place at a negligible rate unless sufficient heat from a spark overcomes their activation energy. However, once they begin to burn, the chemical reactions release enough heat to continue the burning process, supplying the activation energy for surrounding fuel molecules. Like these reactions outside of cells, the activation energy for most cellular reactions is too high for heat energy to overcome at efficient rates. In other words, in order for important cellular reactions to occur at appreciable rates (number of reactions per unit time), their activation energies must be lowered ((Figure)). Scientist refer to this as catalysis. 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 cell’s essential components would disintegrate.
If no activation energy were required to break down sucrose (table sugar), would you be able to store it in a sugar bowl?
Energy comes in many different forms. Objects in motion do physical work, and kinetic energy is the energy of objects in motion. Objects that are not in motion may have the potential to do work, and thus, have potential energy. Molecules also have potential energy because breaking molecular bonds has the potential to release energy. Living cells depend on harvesting potential energy from molecular bonds to perform work. Free energy is a measure of energy that is available to do work. A system’s free energy changes during energy transfers such as chemical reactions, and scientists refer to this change as ∆G.
A reaction’s ∆G can be negative or positive, meaning that the reaction releases energy or consumes energy, respectively. A reaction with a negative ∆G that gives off energy is an exergonic reaction. One with a positive ∆G that requires energy input is an endergonic reaction. Exergonic reactions are spontaneous because their products have less energy than their reactants. Endergonic reactions’ products have a higher energy state than the reactants, and so these are nonspontaneous reactions. However, all reactions (including spontaneous -∆G reactions) require an initial energy input in order to reach the transition state, at which they will proceed. This initial input of energy is the activation energy.
Visual Connection Questions
(Figure) Look at each of the processes, and decide if it is endergonic or exergonic. In each case, does enthalpy increase or decrease, and does entropy increase or decrease?
(Figure) A compost pile decomposing is an exergonic process enthalpy increases (energy is released) and entropy increases (large molecules are broken down into smaller ones). A baby developing from a fertilized egg is an endergonic process enthalpy decreases (energy is absorbed) and entropy decreases. Sand art being destroyed is an exergonic process there is no change in enthalpy, but entropy increases. A ball rolling downhill is an exergonic process enthalpy decreases (energy is released), but there is no change in entropy.
(Figure) If no activation energy were required to break down sucrose (table sugar), would you be able to store it in a sugar bowl?
(Figure) No. We can store chemical energy because of the need to overcome the barrier to its breakdown.
Consider a pendulum swinging. Which type(s) of energy is/are associated with the pendulum in the following instances: i. the moment at which it completes one cycle, just before it begins to fall back towards the other end, ii. the moment that it is in the middle between the two ends, and iii. just before it reaches the end of one cycle (just before instant i.).
- i. potential and kinetic, ii. potential and kinetic, iii. kinetic
- i. potential, ii. potential and kinetic, iii. potential and kinetic
- i. potential, ii. kinetic, iii. potential and kinetic
- i. potential and kinetic, ii. kinetic iii. kinetic
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.
Which of the following is the best way to judge the relative activation energies between two given chemical reactions?
- Compare the ∆G values between the two reactions.
- Compare their reaction rates.
- Compare their ideal environmental conditions.
- Compare the spontaneity between the two reactions.
Critical Thinking Questions
Explain in your own words the difference between a spontaneous reaction and one that occurs instantaneously, and what causes this difference.
A spontaneous reaction is one that has a negative ∆G and thus releases energy. However, a spontaneous reaction need not occur quickly or suddenly like an instantaneous reaction. It may occur over long periods due to a large energy of activation, which prevents the reaction from occurring quickly.
Describe the position of the transition state on a vertical energy scale, from low to high, relative to the position of the reactants and products, for both endergonic and exergonic reactions.
The transition state is always higher in energy than the reactants and the products of a reaction (therefore, above), regardless of whether the reaction is endergonic or exergonic.
German-British medical doctor and biochemist Hans Krebs' 1957 book Energy Transformations in Living Matter (written with Hans Kornberg)  was the first major publication on the thermodynamics of biochemical reactions. In addition, the appendix contained the first-ever published thermodynamic tables, written by Kenneth Burton, to contain equilibrium constants and Gibbs free energy of formations for chemical species, able to calculate biochemical reactions that had not yet occurred.
Non-equilibrium thermodynamics has been applied for explaining how biological organisms can develop from disorder. Ilya Prigogine developed methods for the thermodynamic treatment of such systems. He called these systems dissipative systems, because they are formed and maintained by the dissipative processes that exchange energy between the system and its environment, and because they disappear if that exchange ceases. It may be said that they live in symbiosis with their environment. Energy transformations in biology are dependent primarily on photosynthesis. The total energy captured by photosynthesis in green plants from the solar radiation is about 2 x 10 23 joules of energy per year.  Annual energy captured by photosynthesis in green plants is about 4% of the total sunlight energy that reaches Earth. The energy transformations in biological communities surrounding hydrothermal vents are exceptions they oxidize sulfur, obtaining their energy via chemosynthesis rather than photosynthesis.
The field of biological thermodynamics is focused on principles of chemical thermodynamics in biology and biochemistry. Principles covered include the first law of thermodynamics, the second law of thermodynamics, Gibbs free energy, statistical thermodynamics, reaction kinetics, and on hypotheses of the origin of life. Presently, biological thermodynamics concerns itself with the study of internal biochemical dynamics as: ATP hydrolysis, protein stability, DNA binding, membrane diffusion, enzyme kinetics,  and other such essential energy controlled pathways. In terms of thermodynamics, the amount of energy capable of doing work during a chemical reaction is measured quantitatively by the change in the Gibbs free energy. The physical biologist Alfred Lotka attempted to unify the change in the Gibbs free energy with evolutionary theory.
Energy transformation in biological systems Edit
The sun is the primary source of energy for living organisms. Some living organisms like plants need sunlight directly while other organisms like humans can acquire energy from the sun indirectly.  There is however evidence that some bacteria can thrive in harsh environments like Antarctica as evidence by the blue-green algae beneath thick layers of ice in the lakes. No matter what the type of living species, all living organisms must capture, transduce, store, and use energy to live.
The relationship between the energy of the incoming sunlight and its wavelength λ or frequency ν is given by
where h is the Planck constant (6.63x10 −34 Js) and c is the speed of light (2.998x10 8 m/s). Plants trap this energy from the sunlight and undergo photosynthesis, effectively converting solar energy into chemical energy. To transfer the energy once again, animals will feed on plants and use the energy of digested plant materials to create biological macromolecules.
Thermodynamic Theory of Evolution Edit
The biological evolution may be explained through a thermodynamic theory. The four laws of thermodynamics are used to frame the biological theory behind evolution. The first law of thermodynamics states that energy can not be created or destroyed. No life can create energy but must obtain it through its environment. The second law of thermodynamics states that energy can be transformed and that occurs everyday in lifeforms. As organisms take energy from their environment they can transform it into useful energy. This is the foundation of tropic dynamics.
The general example is that the open system can be defined as any ecosystem that moves toward maximizing the dispersal of energy. All things strive towards maximum entropy production, which in terms of evolution, occurs in changes in DNA to increase biodiversity. Thus, diversity can be linked to the second law of thermodynamics. Diversity can also be argued to be a diffusion process that diffuses toward a dynamic equilibrium to maximize entropy. Therefore, thermodynamics can explain the direction and rate of evolution along with the direction and rate of succession. 
First Law of Thermodynamics Edit
The First Law of Thermodynamics is a statement of the conservation of energy though it can be changed from one form to another, energy can be neither created nor destroyed.  From the first law, a principle called Hess's Law arises. Hess’s Law states that the heat absorbed or evolved in a given reaction must always be constant and independent of the manner in which the reaction takes place. Although some intermediate reactions may be endothermic and others may be exothermic, the total heat exchange is equal to the heat exchange had the process occurred directly. This principle is the basis for the calorimeter, a device used to determine the amount of heat in a chemical reaction. Since all incoming energy enters the body as food and is ultimately oxidized, the total heat production may be estimated by measuring the heat produced by the oxidation of food in a calorimeter. This heat is expressed in kilocalories, which are the common unit of food energy found on nutrition labels. 
Second Law of Thermodynamics Edit
The Second Law of Thermodynamics is concerned primarily with whether or not a given process is possible. The Second Law states that no natural process can occur unless it is accompanied by an increase in the entropy of the universe.  Stated differently, an isolated system will always tend to disorder. Living organisms are often mistakenly believed to defy the Second Law because they are able to increase their level of organization. To correct this misinterpretation, one must refer simply to the definition of systems and boundaries. A living organism is an open system, able to exchange both matter and energy with its environment. For example, a human being takes in food, breaks it down into its components, and then uses those to build up cells, tissues, ligaments, etc. This process increases order in the body, and thus decreases entropy. However, humans also 1) conduct heat to clothing and other objects they are in contact with, 2) generate convection due to differences in body temperature and the environment, 3) radiate heat into space, 4) consume energy-containing substances (i.e., food), and 5) eliminate waste (e.g., carbon dioxide, water, and other components of breath, urine, feces, sweat, etc.). When taking all these processes into account, the total entropy of the greater system (i.e., the human and her/his environment) increases. When the human ceases to live, none of these processes (1-5) take place, and any interruption in the processes (esp. 4 or 5) will quickly lead to morbidity and/or mortality.
Gibbs Free Energy Edit
In biological systems, in general energy and entropy change together. Therefore, it is necessary to be able to define a state function that accounts for these changes simultaneously. This state function is the Gibbs Free Energy, G.
- H is the enthalpy (SI unit: joule)
- T is the temperature (SI unit: kelvin)
- S is the entropy (SI unit: joule per kelvin)
The change in Gibbs Free Energy can be used to determine whether a given chemical reaction can occur spontaneously. If ∆G is negative, the reaction can occur spontaneously. Likewise, if ∆G is positive, the reaction is nonspontaneous.  Chemical reactions can be “coupled” together if they share intermediates. In this case, the overall Gibbs Free Energy change is simply the sum of the ∆G values for each reaction. Therefore, an unfavorable reaction (positive ∆G1) can be driven by a second, highly favorable reaction (negative ∆G2 where the magnitude of ∆G2 > magnitude of ∆G1). For example, the reaction of glucose with fructose to form sucrose has a ∆G value of +5.5 kcal/mole. Therefore, this reaction will not occur spontaneously. The breakdown of ATP to form ADP and inorganic phosphate has a ∆G value of -7.3 kcal/mole. These two reactions can be coupled together, so that glucose binds with ATP to form glucose-1-phosphate and ADP. The glucose-1-phosphate is then able to bond with fructose yielding sucrose and inorganic phosphate. The ∆G value of the coupled reaction is -1.8 kcal/mole, indicating that the reaction will occur spontaneously. This principle of coupling reactions to alter the change in Gibbs Free Energy is the basic principle behind all enzymatic action in biological organisms.