Endergonic and Exergonic Reactions# - Biology

Endergonic and Exergonic Reactions# - Biology

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Endergonic and exergonic reactions

For reactions with ∆G < 0, the products of the reaction have less free energy than the reactants. Reactions that have a negative ∆G are termed exergonic reactions. Understanding which chemical reactions are spontaneous is extremely useful for biologists who are trying to understand whether a reaction is likely to "go" or not.

It is important to note that the term "spontaneous"—in the context of thermodynamics—does NOT imply anything about how fast the reaction proceeds. The change in free energy only describes the difference between beginning and end states, NOT how fast that transition takes place. This is somewhat contrary to the everyday use of the term, which usually carries the implicit understanding that something happens quickly. As an example, the oxidation/rusting of iron is a spontaneous reaction. However, an iron nail exposed to air does not rust instantly—it may take years.

A chemical reaction with a positive ∆G means that the products of the reaction have a higher free energy than the reactants (see the right panel of Figure 1). These chemical reactions are called endergonic reactions, and they are NOT spontaneous. An endergonic reaction will not take place on its own without the transfer of energy into the reaction or increase of entropy somewhere else.

Figure 1. Exergonic and endergonic reactions result in changes in Gibbs free energy. In an exergonic reaction, the free energy of the products is lower than that of the reactants; meanwhile, in an endergonic reaction, the free energy of the products is higher than that of the reactants. Attribution: Marc T. Facciotti (own work)

The building of complex molecules, such as sugars, from simpler ones is an anabolic process and is endergonic. On the other hand, the catabolic process, such as the breaking down of sugar into simpler molecules, is generally exergonic. Like the example of rust above, while the breakdown of biomolecules is generally spontaneous, these reactions don’t necessarily occur instantaneously (quickly). Remember, the terms endergonic and exergonic only refer to the difference in free energy between the products and reactants; they don't tell you about the rate of the reaction (how fast it happens). The issue of rate will be discussed in later sections.

An important concept in the study of metabolism and energy is that of chemical equilibrium. Most chemical reactions are reversible. They can proceed in both directions, often transferring energy into their environment in one direction and transferring energy in from the environment in the other direction. 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 a state of equilibrium is reached. This state of equilibrium is one of the lowest possible free energy states and is a state of maximal entropy. Equilibrium in a chemical reaction is the state in which both reactants and products are present in concentrations that have no further tendency to change with time. Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction. NOTE THIS LAST STATEMENT! Equilibrium means that the relative concentrations of reactants and products are not changing in time, BUT it does NOT mean that there is no interconversion between substrates and products—it just means that when the reactant(s) are converted to product(s) that product(s) are converted to reactant(s) at an equal rate (see Figure 2).

Either a rebalancing of substrate or product concentrations (by adding or removing substrate or product) or a positive change in free energy, typically by the transfer of energy from outside the reaction, is required to move a reaction out of a state of equilibrium. In a living cell, most chemical reactions do not reach a state of equilibrium—this would require that they reach their lowest free energy state. Energy is therefore required to keep biological reactions out of their equilibrium state. In this way, living organisms are in a constant, energy-requiring, uphill battle against equilibrium and entropy.

Figure 2. At equilibrium, do not think of a static, unchanging system. Instead, picture molecules moving in equal amounts from one area to another. Here, at equilibrium, molecules are still moving from left to right and right to left. The net movement however, is equal. There will still be about 15 molecules in each side of this flask once equilibrium is reached. Source:

Metabolic Pathways

Sugar metabolism is a classic example of one of the many cellular processes that use and produce energy. Living things consume sugar as a major energy source due to the high energy stored within their bonds. During photosynthesis, plants use solar energy to convert carbon dioxide gas (CO2) into sugar molecules (like glucose: C6H12O6). Oxygen is produced as a waste product. This reaction is summarized as:

It requires energy input to proceed. During photosynthesis, energy is provided by adenosine triphosphate (ATP), the primary energy currency of all cells. Just as the dollar is used as currency to buy goods, cells use ATP molecules as energy currency to perform work. During cellular respiration, glucose is used as an energy source. It can be summarized by the reverse reaction to photosynthesis. In this reaction, oxygen is consumed and carbon dioxide is released as a waste product. The reaction is summarized as:

Although simplified, both of these reactions involve many steps. A metabolic pathway is very organized series of linked chemical reactions. Two opposite processes are involved. Anabolic pathways require energy input in order to produce large molecules(polymers). Catabolic pathways release energy by breaking down polymers into their smaller molecules(monomers). Consequently, metabolism is composed of synthesis (anabolism) and degradation (catabolism) (Figure 2).

Figure 2. Catabolic pathways are those that generate energy by breaking down larger molecules. Anabolic pathways are those that require energy to synthesize larger molecules. Both types of pathways are required for maintaining the cell’s energy balance.

Chemical Equilibrium

An important concept in the study of 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 (see image below).

Exergonic and endergonic reactions result in changes in Gibbs free energy. Exergonic reactions release energy endergonic reactions require energy to proceed. Image credit: OpenStax Biology

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 a state of equilibrium is reached. This state of equilibrium is one of the lowest possible free energy and a state of maximal entropy.

Energy must be put into the system to push the reactants and products away from a state of equilibrium. 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 work needed 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 chemical equilibrium is never reached. In this way, living organisms are in a constant energy-requiring, uphill battle against equilibrium and entropy. This constant supply of energy ultimately comes from sunlight, which is used to produce nutrients in the process of photosynthesis.

Video Animation Describing Chemical Equilibrium

When molecules collide, chemical reactions can occur. This causes major structural changes akin to getting a new arm on your face. In this TED-Ed video, George Zaidan and Charles Morton playfully imagine chemical systems as busy city streets, and the colliding molecules within them as your average, limb-swapping joes.

[Attributions and Licenses]

This modified article is licensed under a CC BY-NC-SA 4.0 license.

Note that the video(s) in this lesson are provided under a Standard YouTube License.

Examples of endergonic in the following topics:

Free Energy

  • These chemical reactions are called endergonic reactions they are non-spontaneous.
  • An endergonic reaction will not take place on its own without the addition of free energy.
  • Therefore, the chemical reactions involved in anabolic processes are endergonic reactions.
  • Exergonic and endergonic reactions result in changes in Gibbs free energy.
  • Exergonic reactions release energy endergonic reactions require energy to proceed.

Activation Energy

  • Activation energy must be considered when analyzing both endergonic and exergonic reactions.
  • Cells will at times couple an exergonic reaction $(Delta G 0)$, allowing them to proceed.
  • The free energy released from the exergonic reaction is absorbed by the endergonic reaction.
  • Whether the reaction is exergonic (ΔG 0) determines whether the products in the diagram will exist at a lower or higher energy state than the reactants.
  • In this endergonic reaction, activation energy is still required to transform the reactants A + B into the product C.

ATP: Adenosine Triphosphate

  • Cells couple the exergonic reaction of ATP hydrolysis with endergonic reactions to harness the energy within the bonds of ATP.
  • ATP provides the energy for both energy-consuming endergonic reactions and energy-releasing exergonic reactions, which require a small input of activation energy.
  • Cells couple the exergonic reaction of ATP hydrolysis with the endergonic reactions of cellular processes.
  • By donating free energy to the Na+/K+ pump, phosphorylation drives the endergonic reaction.
  • In this example, the exergonic reaction of ATP hydrolysis is coupled with the endergonic reaction of converting glucose for use in the metabolic pathway.

ATP in Metabolism

  • In this way, ATP is a direct link between the limited set of exergonic pathways of glucose catabolism and the multitude of endergonic pathways that power living cells.
  • During an endergonic chemical reaction, ATP forms an intermediate complex with the substrate and enzyme in the reaction.
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Exergonic and Endergonic Reactions - Research Article from World of Biology

Exergonic refers to chemical reactions that proceed spontaneously from reactants to products with the release of energy. Endergonic reactions require energy input to proceed. Although the terms are often used rather loosely, they are precisely defined thermodynamic concepts based on changes in an entity called Gibbs free energy (G) accompanying reactions. Reactions in which -G decreases are exergonic, and those in which -G increases are endergonic. Exergonic reactions often involve the breakdown of organic compounds found in food, whereas endergonic reactions frequently entail synthesis of complicated molecules. Biological metabolism contains many examples of both types, and living organisms have developed elaborate techniques for coupling the two.

Although a negative -G indicates that energy must be added to the system before a reaction will occur, it tells us nothing about the rate at which it will progress. As is often the case, it may go very slowly if substantial activation energy is required to start the reaction. Living organisms have found a way around this problem by forming protein catalysts, called enzymes, that effectively reduce the amount of activation energy needed, and allow the reaction to proceed at a satisfactory rate. Enzymes do not affect the free energy of the reaction, and will not enable reactions to proceed that are not energetically feasible.

By coupling exergonic and endergonic reactions, organisms are able to use the available energy in food they consume to construct complex proteins, lipids, nucleic acids and carbohydrates needed for their growth and development. A well-known example involves coupling the formation of energy-rich adenosine triphosphate (ATP) from adenosine diphosphate (ADP) and phosphate (an endergonic reaction), with the transfer of hydrogen, removed from organic food materials, to oxygen (an exergonic reaction). The process is called oxidative phosphorylation. Energy stored in ATP may be used subsequently when the exergonic conversion of ATP back to ADP and phosphate is coupled with the endergonic synthesis of a needed cellular component.

What you'll learn:

Even though exergonic reactions spontaneous, the reaction will occur at an observable rate. The disproportion of hydrogen peroxide release free energy at a very slow rate if the specific catalysts are not present. The endergonic and exergonic are based on the change of free energy. Exothermic and endothermic reactions depend on the change of enthalpy in a closed system associated with the absorption or release of heat. The exothermic chemical reactions are mostly exothermic due to the breaking of chemical bonds and release of energy. Catabolism refers to the reactions in which the chemical bonds are broken during the metabolism.

Exergonic, Endergonic and Entropy!

If the reaction is spontaneous (Exergonic) does this mean entropy should increase?

1. Oxygen and Hydrogen gases can be reacted together to form water. The entropy of the two gases is higher than the entropy of liquid water, so entropy seems to decrease. However, the reaction is spontaneous. Explain the apparent contradiction.

For any exergonic (spontaneous) reaction, the entropy of the universe will increase. It is very feasible that the entropy of the reacting system will decrease, but still, then entropy of the surroundings will increase more than the decrease in the entropy of the system. Freezing of water at below 0C is a very nice example. Obviously, this is a spontaneous process--you can witness it occurring. As liquid water freezes, entropy of the molecules of water molecules decreases a whole lot. However, in order to freeze liquid water, heat must be removed from the water. That heat increases the entropy of the surrounding air molecules more than the entropy of the water decreases. So, the total entropy of the universe does increase during this spontaneous process.

In this reaction, a whole lot of heat is released when H2 and O2 react to form liquid water. That heat raises the entropy of the molecules all around the water more than the entropy of the H2 and O2 decreases as it forms water.

Exergonic reactions

These reactions are irreversible reactions which occur spontaneously in nature. By spontaneous it means ready or eager to happen with very little external stimuli. Example is combustion of sodium when exposed to oxygen present in the atmosphere. Burning of a log is another example of exergonic reactions. Such reactions liberate more heat and are called as favourable reactions in the field of chemical thermodynamics. The Gibbs free energy is negative under constant temperature and pressure which means that more energy is released rather than absorbed. These are irreversible reactions.

Cellular respiration is a classic example of exergonic reaction. Around 3012 kJ of energy is released when one molecule of glucose is converted to carbon dioxide. This eneegy is utilised by the organisms for other cellular activities. All catabolic reactions i.e. break down of large molecule into smaller molecules is an exergonic reaction. For example – carbohydrate, fat and protein breakdown released energy for the living organisms to do work.

Some exergonic reactions do not occur spontaneously and require a small input of energy to start the reaction. This input of energy is called activation energy. Once the activation energy requirement is fulfilled by an outside source, the reaction proceeds to break bonds and form new bonds and energy is released as the reaction takes place. This results in a net gain in energy in the surrounding system and a net loss in energy from the reaction system.

What is Endergonic

Endergonic is a type of reaction that has a positive Gibbs free energy. The Gibbs free energy is a thermodynamic potential that is used to predict whether a chemical reaction is spontaneous or non-spontaneous. A negative Gibbs free energy indicates a spontaneous reaction. In case of endergonic reactions, the Gibbs free energy is a positive value, which indicates it is a no-spontaneous reaction. Non-spontaneous reactions can also be named as unfavorable reactions.

The Gibbs free energy of endergonic reaction is a positive value when calculated using the following thermodynamic relationship.


Where, ΔG is the Gibbs free energy

T is the temperature of the system

Figure 1: Energy Diagram for an Endergonic Reaction

In a non-spontaneous reaction, energy should be provided from outside for the progression of the reaction. Then, the energy of the products gets a higher value than that of the energy of the reactants. Due to that reason, the change in enthalpy is a positive value (the change in the enthalpy is the difference between the enthalpies of products and reactants). Since new products are formed, the entropy of the system is decreased. Then, according to the above equation, the ΔG is a positive value. Endergonic reactions include endothermic reactions.

Plant Life

The primary source of energy for life on the earth is the sun,which is the energy source for photosynthesis: the biological process that transforms radiant energy into chemical energy. Chemical energy is stored in biological molecules, which can then be used as the fuel to provide an organism’s energy needs.

Such biological molecules include sugars (or carbohydrates), proteins, and lipids (or fats). In the reactions of metabolism, many types of molecules are synthesized (anabolism), and many are broken down (catabolism). Changes in energy content occur in all these reactions.

Bioenergetics is the science that studies the description of the basic mechanisms that govern the transformation and use of energy by organisms. A basic tenet of bioenergetics is that no chemical reaction can be 100 percent energy-efficient. In other words, in all reactions there is some transfer of energy, but some of it is always lost in the form of heat.

The energy (often measured in calories) contained in the molecular structure of a compound is called Gibbs free energy (after Josiah Willard Gibbs, 1839-1903, who founded the discipline of physical science) and is the energy available to perform work.

The difference between the free energy of the products and the free energy of the reactants in a chemical reaction is called the change in free energy and is fundamental in determining if a reaction can occur spontaneously. If the change in free energy is negative, energy is released, and the free energy content is less in the products than in the reactants.

Such reactions are considered exergonic. On the other hand, if the change in free energy is positive, the reaction is considered endergonic and is non-spontaneous (that is, endergonic reactions require a source of energy to enable them to occur).

Many cellular reactions are endergonic and can not occur spontaneously. Nevertheless, cells can facilitate endergonic reactions using the energy released from other exergonic reactions, a process called energy coupling.

As an example, consider a common endergonic reaction in plants in which glucose and fructose are joined together to make sucrose. To enable this reaction to take place, it is coupled with a series of other exergonic reactions as follows:

glucose + adenosine triphosphate (ATP) → glucose-p + ADP

fructose + ATP → fructose-p + adenosine diphosphate (ADP)

glucose-p + fructose-p → sucrose + 2 Pi(inorganic phosphate)

Therefore, although producing sucrose from glucose and fructose is an endergonic reaction, all three of the foregoing reactions are exergonic. This is representative of the way cells facilitate endergonic reactions.

The principal molecule involved in providing the energy for endergonic cellular reactions to take place is adenosine triphosphate, or ATP, the same molecule used in the example above.

ATP is typically produced by joining an inorganic phosphate to adenosine diphosphate (ADP), which is an endergonic reaction. This, too, represents a characteristic of chemical reactions: If a reaction is exergonic in one direction, it will be endergonic in the opposite direction.

Thus, the breakdown of ATP is exergonic, while the production of ATP is endergonic. The energy for production of most of the ATP in plant cells comes from the light reactions of photosynthesis and the electron transport system in the mitochondria.

The enigma is why ATP, and not any other molecule, is used. Although no complete justification is available, there are several points that support its significance.

First, there is the high stability of the ATPmolecule at the physiological pH (around 7.4) toward hydrolysis and decomposition in the absence of an enzyme catalyst. This stability allows ATP to be stored in the cell until needed. Second, ATP is one of the molecules (a nucleotide) that is used in synthesis ofDNA.

Finally, the magnitude of the change in free energy involved in the ATP-ADP transformation is of an amount useful for driving many of the endergonic reactions in the cell. As a result, it can play the role of an intermediate quite easily.

Watch the video: endergonic and exergonic reactions (August 2022).