C. Change in Free Energy G - Biology

C. Change in Free Energy G - Biology

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The total ΔG can be expressed as the sum of the two contributions showing the effects of the intrinsic stability (Keq) and concentration:

ΔG = ΔGstab + ΔGconc

ΔG = ΔGo + RTlnQrx = ΔGo + RTln ([P][Q])/([A][B])

for the reaction A + B <=> P + Q, where ΔGoreflects the contribution from the relative intrinsic stability of reactants and products) and RTlnQrx reflects the contribution from the relative concentrations of reactants and products (which has nothing to do with stability). Qrx is the reaction quotient which for the reaction A + B <=> P + Q is given by:

Qrx = ([P][Q])/([A][B])

Meaning of ΔG

Remember that ΔG is the "driving" force for a reaction, analogous to the difference in potential energy for a ball on a hill. Go back to that analogy. if the ball starts at the top of the hill, does it roll down hill? Of course. It goes from high potential energy to low potential energy. The reaction can be written as: Balltop --> Ballbottom for which the change in potential energy, ΔPE = PEbottom -PEtop< 0. If the ball starts at the bottom, will it go to the top? Obviously not. For that reaction, Ballbottom --> Balltop, ΔPE > 0. If the top of the hill was at the same height at the bottom of the hill (obviously an absurd situation), the ball would not move. It would effectively be at equilibrium, a state of no change. For this reaction, Balltop --> Ballbottom, the ΔPE = 0. As the ball starts rolling down the hill, its potential energy gets closer to the potential it would have at the bottom. Hence the ΔPE changes from negative to more and more positive until it gets to the bottom at which case the ΔPE = 0 and movement ceases. If the ΔPE is not 0, the ball will move until the ΔPE = 0.

Likewise, for a chemical reaction that favors products, ΔG < 0. The system is not at equilibrium and the reaction will go in the direction of products. As the reaction proceeds, products build up, and there is less of a driving force for reactants to go to products (LeChatilier's Principle), so the ΔG becomes more a more positive until the ΔG = 0 and the reaction is at equilibrium. A reaction that has a ΔG > 0 is likewise not at equilibrium so it will go in the appropriate direction until equilibrium is reached. Hence for the reaction A + B <==> P + Q,

  • if ΔG < 0, the reaction goes toward products P and Q
  • if ΔG = 0, the reaction is at equilibrium and no further change occurs in the concentration of reactants and products.
  • if ΔG > 0, the reaction goes toward reactants A and B.

We can not measure easily the actual free energy G of reactants or products, but we can measure ΔG readily. These points are illustrated in the graph below of ΔG vs time for the hypothetical reaction A + B <==> P + Q. (Also notice the two insert graphs - in blue and red - which show, in analogy to the ball on the hill graphs, the values of ΔG at the two points where the perturbation to the equilibrium were made.)

Notice the ΔG is constantly changing until the system reaches equilibrium. Initially the equilibrium is perturbed so that the system is not in equilibrium (shown in blue). The perturbation was such that the products are favored. After equilibrium was reached, the system was perturbed again, this time in a fashion to favor the reverse reaction. Notice in this case the ΔG for the reaction as written: A + B <==> P + Q is positive - i.e. it is not in equilibrium. Therefore the reaction (as written) goes backwards to products. It is important to realize that the reported ΔG is for the reaction as written.

Now let's apply ΔG = ΔGo + RTln Q = ΔGo + RTln ([P][Q])/([A][B]) to two reactions we discussed above:

  • HCl(aq) + H2O(l) <==>H3O+(aq) + Cl-(aq)
  • CH3CO2H(aq) + H2O(l) <==> H3O+(aq) + CH3CO2-(aq)

Assume that at time t=0, 0.1 mole of HCl and CH3CO2H were added to two different beakers. At this point the forward reaction are favored, but obviously to different extents. The RTln Q would be identical for both acids, since each reactant is present at 0.1 M, but no products yet exist. However, the ΔGois negative for HCl and positive for acetic acid since HCl is a strong acid. Hence at t=0, ΔG for the HCl reaction is much more negative than for acetic acid. This is summarized in table below. The direction of the arrow shows if products (-->) or reactants (<---) are favored. The size of the arrow shows very approximately to what extent the ΔG term is favored

Reaction at t=0


RTln Q


HCl(aq) + H2O(l)




CH3CO2H(aq) + H2O(l)




Now when equilibrium is reached, no net change occurs in the concentration of reactants and products, and ΔG = 0. In the case of HCl, there is just an infinitesimal amount of HCl left, and 0.1 M of each product, so concentration favors HCl formation. However, the intrinsic relative stability of reactants and products still favors products. In the case of acetic acid, most of the acetic acid remains (0.099 M) with little product (0.001 M) so concentration favors product. However, the intrinsic relative stability of reactants and products still favors reactants. This is summarized in table below.

Reaction at equlib.


RTln Q


HCl(aq) + H2O(l)



favors neither, = 0

CH3CO2H(aq) + H2O(l)



favors neither, = 0

Compare the two tables above (one at time t= 0 and the other at equilibrium). Notice:

  • ΔGonever changes, since it has nothing to do with concentration.
  • Only RTln Q changes during the course of a reaction, until equilibrium is achieved.

Meaning of ΔGo

To get a better meaning of the significance of ΔGo, let's consider the following equations under two different conditions:

ΔG = ΔGo + RTln Q = ΔGo + RTln ([P][Q])/([A][B])

Condition I: Reaction at equilibrium, DG = 0

The equation reduces to: ΔGo = - RTln ([P]eq[Q]eq)/([A]eq[B]eq) or ΔGo = - RTln Keq = - 2.303RTlog Keq

This supports our idea that DGo is independent of concentration since Keq should also be independent of concentration.

Condition II: Concentration of all reactants and products is 1 M (standard state, assuming solution reaction)

The equation reduces to: ΔG = ΔGo + RTln ([1][1])/([1][1]) = ΔGo + 2.303RTlog 1 = ΔGo

The implies that when all reactants are at this concentration, defined as the standard state (1 M for solutes), the ΔG at that particular moment just happens to be the DGo for the reaction. If one of the reactant or products is H3O+, it would make little biological sense to calculate DGo for the reaction using the standard state of [H3O+] = 1 M, or a pH of -1. Instead, it is assumed the pH = 7, [H3O+] = 10-7 M. A new symbol is used for DGo under these condition, ΔGo'.

H + H --> H2 Does this reaction occur spontaneously? It does. You should remember that individual H atoms are unstable, since they don't have an completed outer shell of electrons - in this case a duet. As they approach, they can interact to form a covalent bond and in the process release energy. The bonded state is a lower energy state than two separated H atoms. This should be clear since energy has to be added to a molecule of H2 to break the bond.


2C8H18(l) + 25O2(g) --> 16CO2(g) + 18H2O(g). To carry out this reaction, every C-C, C-H and O-O bond in the reactants must be broken (which requires an input of energy) but a lots of energy is released on formation of C-O and H-O covalent bonds in the products. Is more energy needed to break the bonds in the reactants or is more energy released on formation of bonds in the product? The answer should be clear. The products must be at a lower energy than the reactants since huge amounts of heat and light energy are released on combustion of gasoline and other hydrocarbons.

These reactions suggest that energy must be released for a reaction to proceed to any extent in a given direction.

Now consider, however, the following reaction:

Ba(OH)2.8H2O(s) + 2NH4SCN(s) --> 10H2O(l) + 2NH3(g) + Ba(SCN)2(aq+s)

When these two solids are mixed, and stirred, a reaction clearly takes places, as evidenced by the formation of a liquid (water) and the smell of ammonia. What is surprising is that heat is not released into the surroundings in this reaction. Rather heat was absorbed from the surroundings turning the beaker so cold that it freezes to a piece of wood (with a layer of water added to the wood) on which it was placed. This reaction seems to violate our idea that a reaction proceeds in a direction in which heat is liberated. Reactions which liberate heat and raise the temperature of the surroundings are called exothermic reactions. Reactions which absorb heat from the surroundings and hence lower the temperature of the surroundings are endothermic reactions.

C. Change in Free Energy G - Biology

Free energy, called Gibbs free energy (G), is usable energy or energy that is available to do work.

Learning Objectives

Discuss the concept of free energy.

Key Takeaways

Key Points

  • Every chemical reaction involves a change in free energy, called delta G (∆G).
  • To calculate ∆G, subtract the amount of energy lost to entropy (∆S) from the total energy change of the system this total energy change in the system is called enthalpy (∆H ): ΔG=ΔH−TΔS.
  • Endergonic reactions require an input of energy the ∆G for that reaction will be a positive value.
  • Exergonic reactions release free energy the ∆G for that reaction will be a negative value.

Key Terms

  • exergonic reaction: A chemical reaction where the change in the Gibbs free energy is negative, indicating a spontaneous reaction
  • endergonic reaction: A chemical reaction in which the standard change in free energy is positive, and energy is absorbed
  • Gibbs free energy: The difference between the enthalpy of a system and the product of its entropy and absolute temperature

Free Energy

Since chemical reactions release energy when energy-storing bonds are broken, how is the energy associated with 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 quantitate these energy transfers. Free energy is called Gibbs free energy (G) after Josiah Willard Gibbs, the scientist who developed the measurement. 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, resulting in entropy. Gibbs free energy specifically refers to the energy associated with a chemical reaction that is available after accounting for entropy. In other words, Gibbs free energy is usable energy or energy that is available to do work.

Calculating ∆G

Every chemical reaction involves a change in free energy, called delta G (∆G). The change in free energy can be calculated for any system that undergoes a change, such as a chemical reaction. To calculate ∆G, subtract the amount of energy lost to entropy (denoted as ∆S) from the total energy change of the system. This total energy change in the system is called enthalpy and is denoted as ∆H. The formula for calculating ∆G is as follows, where the symbol T refers to absolute temperature in Kelvin (degrees Celsius + 273): G=ΔH−TΔS.

The standard free energy change of a chemical reaction is expressed 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. Standard pH, temperature, and pressure conditions are generally calculated at pH 7.0 in biological systems, 25 degrees Celsius, and 100 kilopascals (1 atm pressure), respectively. It is important to note that cellular conditions vary considerably from these standard conditions therefore, standard calculated ∆G values for biological reactions will be different inside the cell.

Endergonic and Exergonic Reactions

If energy is released 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 products of the reaction have less free energy than the reactants because they gave off some free energy during the reaction. Reactions that have a negative ∆G and, consequently, release free energy, are called exergonic reactions. Exergonic means energy is exiting the system. These reactions are also referred to as spontaneous reactions because they can occur without the addition of 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. An important distinction must be drawn 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. The rusting of iron is an example of a spontaneous reaction that occurs slowly, little by little, over time.

If a chemical reaction requires an input of energy 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, the products of these reactions can be thought of as energy-storing molecules. 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.

Exergonic and Endergonic Reactions: Exergonic and endergonic reactions result in changes in Gibbs free energy. Exergonic reactions release energy endergonic reactions require energy to proceed.

Free Energy and Biological Processes

In a living cell, chemical reactions are constantly moving towards equilibrium, but never reach it. 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.

When complex molecules, such as starches, are built from simpler molecules, such as sugars, the anabolic process requires energy. Therefore, the chemical reactions involved in anabolic processes are endergonic reactions. On the other hand, the catabolic process of breaking sugar down into simpler molecules releases energy in a series of exergonic reactions. As in the example of rust above, the breakdown of sugar involves spontaneous reactions, but these reactions don’t occur instantaneously. 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.

Endergonic and Exergonic Processes: Shown are some examples of endergonic processes (ones that require energy) and exergonic processes (ones that release energy). These include (a) a compost pile decomposing, (b) a chick hatching from a fertilized egg, (c) sand art being destroyed, and (d) a ball rolling down a hill.

Equilibrium Constant of a Reaction and Free Energy Change

In this article we will discuss about the equilibrium constant of a reaction and free energy charge.

Any chemical reaction will reach an equilibrium after a sufficient time.

A reversible reaction may be written:

The velocity of the reaction which will produce C and D is written:

and that of the reaction producing A and B would be:

k1 and k2 are the velocity constants of the reaction.

When equilibrium is reached, velocities v1 and v2 are evidently equal

The concentrations indicated in this formulation are those reached at equi­librium.

The capacity of a molecule to react is characterized by a parameter ex­pressed in kCal/mole and called free energy. The free energy change in a reaction of the type A + B ⇋ C + D when one mole of A and one mole of B give one mole of C and one mole of D, while their respective concentrations [A], [B], [C], and [D] are maintained constant, is written:

This change therefore depends on two parameters. First, a constant char­acteristic of the reaction (∆G0), the standard free energy change, which is the change in free energy in conditions called standard: temperature, 298°K pressure, one atmosphere pH = 0 and [A], [B], [C], [D], maintained at 1M. We have then ∆G = ∆G0.

∆G then depends on the concentration of the products and reagents. It is important to note that, theoretically, there are always concentrations of reagents and products such that ∆G < 0. In this case, the reaction is called exergonic and is always accompanied by a decrease in free energy. It can then yield energy.

For example the reaction:

has a ∆G of – 686 kcal/mole of glucose, which means that in standard condi­tions, the oxidation of one mole of glucose will yield 686 kilocalories of free energy. This energy may be dissipated as heat or converted into mechanical energy (muscle contraction), electrical energy (transmission of nerve impulse or formation of ion gradients), or in certain cases, even radiant energy. It can also be conserved as chemical energy in molecules, the most important of which is ATP.

Lastly, if the reaction is in equilibrium, ∆G = 0 and one has the following equation:

There is therefore, a direct relationship between the standard free energy-change of the reaction and its equilibrium constant.

Every chemical reaction involves a change in free energy, called delta G (∆G). The change in free energy can be calculated 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 total energy change of the system. This total energy change in the system is called enthalpy and is denoted as ∆H . The formula for calculating ∆G is as follows, where the symbol T refers to absolute temperature in Kelvin (degrees Celsius + 273):

The standard free energy change of a chemical reaction is expressed 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. Standard pH, temperature, and pressure conditions are generally calculated at pH 7.0 in biological systems, 25 degrees Celsius, and 100 kilopascals (1 atm pressure), respectively. It is important to 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.

Basic Principles of Energy Conservation

In this article some important principles are undertaken so that we can well understand the various mechanisms of energy conservation.

Free Energy:

In microbiology, energy is measured in units of kilojoules (kJ), a measure of heat energy. Chemical reactions are accompanied by changes in energy. Although in any chemical reaction some energy is lost as heat, in microbiology the interest is in free energy (abbreviated G), which is defined as the energy released that is available to do useful work.

The change in free energy during a reaction is expressed as ∆G 0, , where the symbol A should be read “change in”. The “o’ and ” ‘ “(prime) mean that the free energy value was obtained under “standard’ conditions: pH 7, 25°C, all reactants and products initially at 1M concentration.

the ∆G 0, is negative, the reaction will proceed with the release of free energy, energy that the cell may be able to conserve in the form of adinosine triphosphate (ATP). Such energy yielding reactions are called exergonic. However, if ∆G 0, is positive the reaction requires energy in order to proceed, such reactions are called endergonic. Thus, from the standpoint of the microbial cell, exergonic reactions yield energy, while endergonic reactions require energy.

Free Energy of Formation:

Free energy of formation (abbreviated G O F) is the energy yielded or energy required for the formation of a given molecule from its constituent elements. By convention, the free energy of formation (G O F) of the elements (for example, C, H2, N2) is zero.

If the formation of a compound from elements proceeds exergonically, then the free energy of formation of the compound is negative (energy is released), whereas if the reaction is endergonic (energy is required), then the free energy of formation of the compound is positive.

The values of free energy of formation are always in kilojules/molecule (kJ/mol). Using free energies of formation, it is possible to calculate the change in the free energy taking place in a given reaction. For a simple reaction such as A + B → C + D. the change in free energy (∆G 0, ) can be calculated by subtracting the sum of the free energies of formation of the reactants (in this case A and B) from that of the products (in this case C and D).

Change in free energy (∆G 0, ) of

A + B → C + D = G O F [C + D] – G O F [ A + B]

Table 23.1 shows the free energies of formation for some of the compounds of biological interest. For most of the compounds taken in this table the free energy of formation (ΔG O F) is negative, reflecting the fact that compounds tend to form spontaneously from elements. However, the positive free energy of formation for nitrous oxide (N2O + 104.2 kJ/mol) tells us that this molecule does not form spontaneously, but rather decomposes to nitrogen and oxygen.

Oxidation-Reduction (Redox) Reactions:

The free energy in living organisms is conserved involving oxidation-reduction (redox) reactions. Oxidation- reduction (redox) reactions are those in which electrons are donated by an electron donor (oxidation) and are accepted by an electron acceptor (reduction).

The electron donor is called the reducing agent or reductant, whereas the electron acceptor as oxidising agent or oxidant. By convention, such a reaction is written with the reductant to the right of the oxidant and the number (n) of electrons (e – ) donated.

Electron Donors and Electron Acceptors:

Oxidation-reduction reactions, as stated earlier, involve electrons being donated by an electron donor and being accepted by an electron acceptor. However, the electrons released by the electron donor cannot exist free in solution they must be subsequently accepted by an electron acceptor and become the part of it. This is the reason why for any oxidation to occur, a subsequent reduction must also occur.

For example, hydrogen gas (H2) can release electrons and hydrogen ions (protons) and become oxidized:

The above reaction is only a half reaction and needs subsequently the second half reaction to complete.

In second half reaction there can be the reduction of many different substances including O2:

The second half reaction (reduction reaction), when coupled to the first half reaction (oxidation reaction), yields the following overall balanced reaction:

In the above overall balanced reaction, one refers to the H2 oxidised (i.e., electron donor) and O2 reduced (i.e., electron acceptor). The overall view of the formation of H2O from the electron donor H2 and the electron acceptor O2 is shown in Fig. 23.1.

Reduction Potentials:

Reduction potentials (E0’) is the expression of the tendency of substances to become oxidised or to become reduced the substances vary in the tendency to become oxidised or to become reduced.

The reduction potential is measured electrically in units of volts or millivolts in reference to a standard substance the reduction potential of hydrogen (H2) at pH 7 is -0.42 volts or -0.420 millivolts. pH 7 is used because it refers to neutrality and the cytoplasm of most cells is neutral or nearly so.

Redox Couples:

Most molecules can be either electron donors or electrons acceptors under different conditions, depending on the substances with which they react. The same atom on each side of the arrow in the half reactions can be thought of as representing a redox couple. When writing a redox couple, the oxidised form is always placed on the left.

In constructing complete oxidation-reduction reactions from their constituent half reactions, it is simplest to remember that the reduced substance of a redox couple whose reduction potential is more negative donates electrons to the oxidized substance of a redox couple whose reduction potential is more positive.

Thus, in a redox couple 2H + / H2, which has a reduction potential of -0.42 volts, H2has a great tendency to donate electrons. On the other hand, in the redox couple ½ O2/H2O, which has a potential of +0.82 volts, H2O has a very slight tendency to donate electrons, but O2 has a great tendency to accept electrons.

It follows then that in a reaction of H2 and O2, H2 will be the electron donor and become oxidised, and O2 will be the electron acceptor and become reduced (Fig. 23.1).

Electron Tower:

Electron toner is an imaginary vertical tower that represents the range of reduction potentials for redox couples from the most negative at the top to the most positive at the bottom (Fig. 23.2).

The reduced substance in the redox pair at the top of the tower possesses the greatest tendency to donate electrons, whereas the oxidised substance in the couple at the bottom of the tower has the greatest tendency to accept electrons. As electrons from the electron donor at the top of the tower fall, they can be “caught” by acceptors at various levels of the tower.

The farther the electrons drop from a donor before they are “caught” by an acceptor, the greater the amount of energy released. O2, at the bottom of the tower, is the most favourable electron acceptor used by organisms. In the middle of the electron tower, redox couples can act as either electron donors or electron acceptors.

For instance, under conditions where oxygen is absent (called anoxic) in the presence of H2, fumarate can be electron acceptor (producing succinate), and under other conditions where oxygen is present (called aerobic) in the absence of H2, succinate can be an electron donor (producing fumarate).

Electron Carriers:

The transfer of electrons in an oxidation-reduction reaction from donor to acceptor in a cell normally involves one or more intermediates that are called electron carriers (or carriers). In such conditions the initial electron donor is called primary electron donor, whereas the final electron acceptor as the terminal electron acceptor.

The net change in the free energy (∆G 0, ) of the complete reaction sequence is determined by the “difference” in the reduction potentials (E0’) between the primary electron donor and the terminal electron acceptor.

Electron carriers can be divided into two general groups:

(1) Freely diffusible and (2) firmly attached (fixed) to enzymes in the cytoplasmic membrane. Freely diffusible carriers include the coenzymes nicotinamide adenine dinucleotide (NAD + ) and NAD-phosphate (NADP + ) , whereas membrane-associated electron carriers include NADH dehydrogenases, flavoproteins that contain flavin mononucleotide (FMN) or flavin-adenine dinucleotide (FAD), cytochromes, nonheme iron-sulphur (Fe/S) proteins (ferrodoxin) and quinones.

Nicotinamide Adenine Dinucleotide (NAD + ) and NAD-phosphate (NADP + ):

These are the coenzymes that act as freely diffusible electron carriers and transport electrons between two different locations. The nicotinamide ring of NAD + and NADP + (Fig. 23.3) accepts two electrons and one proton from a donor, while a second proton is released. The reduction potential of the redox couple NAD + /NADH (or NADP + /NADPH) is -0.32 volt, which places it fairly high on the electron tower, i.e., NADH (or NADPH) is a good electron donor.

However, although the NAD and NADP + couples possess the same reduction potentials, they generally function in different capacities in the cell. NAD + /NADH is directly involved in energy generating (catabolic) reactions, whereas NADP + /NADPH is involved primarily in biosynthetic (anabolic) reactions.

NADH dehydrogenases:

NADH dehydrogenases are proteins bound to the inside surface of the cell membrane. They accept hydrogen atoms from NADH generated in various cellular reactions and pass the hydrogen atoms to flavoproteins.


Flavoproteins are proteins possessing a derivative of riboflavin. Flavoproteins accept hydrogen atoms and donate electrons. Two flavoproteins are commonly found in cells—flavin mononucleotide (FMN) and flavin- adenine dinucleotide (FAD) (Fig. 23.4).

Flavin mononucleotide (FMN) is bonded to ribose and adenine through a second phosphate. However, these two flavoproteins bear two electrons and two protons (two hydrogen atoms) on their complex ring system. Riboflavin, also called vitamin B2, is a required growth factor for some organisms.


Cytochromes are proteins with iron-containing porphyrin ring (Fig. 23.5) also called heme. The cytochromes undergo oxidation and reduction through loss or gain of a single electron by the iron atom centrally placed in the porphyrin ring of the cytochrome:

Cytochrome – Fe 2+ ⇋ Cytochrome – Fe 3+ + e –

Cytochromes do not carry hydrogen atoms (protons). Several different cytochromes (cyt b, cyt c, etc.) are a prominent part of respiratory electron transport chains.

Nonheme Iron-Sulphur (Fe/S) Proteins:

Some iron-sulphur (Fe/S) containing electron carrying membrane-associated proteins lack a heme group and are called nonheme iron-sulphur (Fe/S) proteins. Various arrangements of iron and sulphur have been found in different nonheme iron-sulphur proteins, but Fe2S2 and Fe4S4 clusters are the most common.

The iron atoms are bonded to free sulphur and to the protein via sulphur atoms from cysteine residues (Fig. 23.6). Ferredoxin is a common iron-sulphur protein of Fe2S2 configuration occurring in biological systems.

This electron carrier is active in photosynthetic electron transport and several other electron transport processes. Since the reduction potentials of iron-sulphur proteins vary over a wide range depending on the number of iron and sulphur atoms and their attachment-pattern to protein, different iron-sulphur proteins function at different points in an electron transport process. Like cytochromes, these proteins also carry electrons only, not hydrogen atoms.


Quinones are membrane-associated highly hydrophobic non-protein-containing molecules that act as electron carriers in electron transport processes. Some quinones occurring in bacteria are related to vitamin K, a growth factor for higher animals.

Like flavoproteins, quinones serve as proton (hydrogen atom) acceptors and electron donors. Coenzyme Q (CoQ) or ubiquinone is a quinone that carries electrons and protons (hydrogen atoms) in many respiratory electron transport processes (Fig. 23.7).

High-Energy Compounds:

Energy released as a result of oxidation-reduction reactions must be conserved so that it can be utilized wherever and whenever required in the cellular functions. Conservation of energy in living organisms is made in high-energy phosphate bonds of high-energy phosphate compounds (e.g., phosphoenolpyruvate, 1,3- bisphosphaglycerate, ATP, ADP, etc.) and thio-ester bonds of the derivatives of coenzyme A (e.g., acetyl- CoA). We take only ATP and coenzyme A (CoA) derivatives (e.g., acetyl-CoA) for further consideration here.

Adenosine Triphosphate (ATP):

The most important high-energy phosphate compound in living organisms is adenosine triphosphate (ATP), a practical form of major energy-currency the cells possess to carry out their work. ATP consists of the ribonucleoside adenosine to which three phosphate molecules are bonded in series (Fig. 23.8).

Out of the three phosphate bonds of ATP, as is apparent in the Fig. 23.8, two are high-energy anhydride bonds having high free energies of hydrolysis, whereas one is low-energy ester bond.

When ATP breaks down to adenosine diphosphate (ADP) and orthrophosphate (Pi) as a result of the hydrolysis of high energy anhydride bond, the free energy is made available to drive biosynthetic reactions and other aspects of cell function through carefully regulated processes in which the energy released from ATP hydrolysis is coupled to energy-requiring reactions. Later, energy from photosynthesis, aerobic respiration, anaerobic respiration, and fermentation is used to resynthesise ATP from ADP and Pi.

Coenzyme A (CoA) Derivatives (Acetyl-CoA):

Coenzyme A (CoA) derivatives are certain other high-energy compounds that are produced in the cell and can conserve the energy released in oxidation-reduction reactions. These derivatives [acetyl-CoA (Fig. 23.9) is just one of many CoA derivatives] possess thio-ester (sulphoanhydride) bonds instead of phosphoanhydride bonds that occur in high energy phosphate compounds (e.g., ATP). CoA derivatives yield sufficient free energy on hydrolysis, which is used to drive the synthesis of a high energy phosphate bond in energy metabolism and biosynthesis of fatty acids.

For instance, in the reaction:

Acetyl-CoA + H2O + ADP + P → Acetate – + HS-CoA + ATP + H +

the energy released in the hydrolysis of coenzyme A is conserved in the synthesis of ATP. CoA derivatives play very important part in the energy conservation of anaerobic microorganisms, especially in those whose energy metabolism involves fermentation.

Options for Energy Conservation:

Metabolism is the sum total of all biochemical reactions that take place in the cell with the involvement of flow of energy and the participation of variety of enzymes and proteins. Metabolism, in fact, represents the chemistry of life and can be divided into two major parts: catabolism and anabolism.

Catabolism (Gk. cata = down, ballein = to through) represents the breakdown of more complex chemicals into smaller, simpler molecules resulting in the release of energy. Some part of this released energy is trapped and made available for cellular functions while the rest is released as heat. In anabolism (Gk. ana = up, ballein = to through) the similar molecules are used in the synthesis of complex molecules with energy utilization.

Enzymes required for metabolic activities are synthesized in the cell, whereas energy is obtained from one of the three sources (Fig. 23.10):

(i) Chemolithotrophic microbes carry out oxidation of inorganic chemicals that releases energy,

(ii) chemoorganotrophic microorganisms oxidize organic molecules to liberate energy, and

(iii) phototrophic microorganisms trap radiant energy of sun by the process of photosynthesis.

The chemotrophic macroorganisms (both chemolithotrophic and chemoorganotrophic), the microorganisms that use chemicals as electron donors in their energy metabolism, have adopted two catabolic mechanisms for energy conservation respiration and fermentation. In respiration, the energy is conserved by the process oxydative phosphorylation with the involvement of molecular oxygen or some other externally derived electron-acceptor.

Respiration, however, is of two different types, namely, aerobic and anacrobic. In aerobic respiration, the final electron acceptor is oxygen whereas the electron-acceptor in anaerobic respiration is more often inorganic (e.g., NO3 – , SO4 2- , CO2, Fe 3+ , SeO4 2- , and many others), though organic electron-acceptors such as fumaric acid may also be used.

In fermentation, the energy is produced by substrate-level-phosphorylation in which ATP is synthesized as a result of the oxidation of an organic compound without involvement of any usable external electron-acceptor. Phototrophic microorganisms employ anabolic mechanism and trap light energy of sun during photosynthesis (synthesis of complex molecule using simpler molecules) by the process photophosphorylation.

Spontaneous Reactions

A spontaneous reaction is a reaction that favors the formation of products at the conditions under which the reaction is occurring. A roaring bonfire (see figure below) is an example of a spontaneous reaction. A fire is exothermic, which means a decrease in the energy of the system as energy is released to the surroundings as heat. The products of a fire are composed mostly of gases such as carbon dioxide and water vapor, so the entropy of the system increases during most combustion reactions. This combination of a decrease in energy and an increase in entropy means that combustion reactions occur spontaneously.

Figure (PageIndex<1>): Combustion reactions, such as this fire, are spontaneous reactions. Once the reaction begins, it continues on its own until one of the reactants (fuel or oxygen) is gone.

A nonspontaneous reaction is a reaction that does not favor the formation of products at the given set of conditions. In order for a reaction to be nonspontaneous, one or both of the driving forces must favor the reactants over the products. In other words, the reaction is endothermic, is accompanied by a decrease in entropy, or both. Out atmosphere is composed primarily of a mixture of nitrogen and oxygen gases. One could write an equation showing these gases undergoing a chemical reaction to form nitrogen monoxide.

[ce left( g ight) + ce left( g ight) ightarrow 2 ce left( g ight)]

Fortunately, this reaction is nonspontaneous at normal temperatures and pressures. It is a highly endothermic reaction with a slightly positive entropy change (left( Delta S ight)). However, nitrogen monoxide is capable of being produced at very high temperatures, and this reaction has been observed to occur as a result of lightning strikes.

One must be careful not to confuse the term spontaneous with the notion that a reaction occurs rapidly. A spontaneous reaction is one in which product formation is favored, even if the reaction is extremely slow. You do not have to worry about a piece of paper on your desk suddenly bursting into flames, although its combustion is a spontaneous reaction. What is missing is the required activation energy to get the reaction started. If the paper were to be heated to a high enough temperature, it would begin to burn, at which point the reaction would proceed spontaneously until completion.

In a reversible reaction, one reaction direction may be favored over the other. Carbonic acid is present in carbonated beverages. It decomposes spontaneously to carbon dioxide and water according to the following reaction.

[ce left( aq ight) ightleftharpoons ce left( g ight) + ce left( l ight)]

If you were to start with pure carbonic acid in water and allow the system to come to equilibrium, more than (99\%) of the carbonic acid would be converted into carbon dioxide and water. The forward reaction is spontaneous because the products of the forward reaction are favored at equilibrium. In the reverse reaction, carbon dioxide and water are the reactants, and carbonic acid is the product. When carbon dioxide is bubbled into water (see figure below), less than (1\%) is converted to carbonic acid when the reaction reaches equilibrium. The reverse of the above reaction is not spontaneous. This illustrates another important point about spontaneity. Just because a reaction is not spontaneous does not mean that it does not occur at all. Rather, it means that the reactants will be favored over the products at equilibrium, even though some products may indeed form.

Figure (PageIndex<2>): A home soda making machine is shown with a bottle of water and a (ce) cartridge. When the water is carbonated, only a small amount of carbonic acid is formed because the reaction is nonspontaneous. (Public Domain Baruchlanda)


Figure 6. Enzymes lower the activation energy of the reaction but do not change the free energy of the reaction.

A substance that helps a chemical reaction to occur is called a catalyst, and the molecules that catalyze biochemical reactions are called enzymes. Most enzymes are proteins and perform the critical task of lowering the activation energies of chemical reactions inside the cell. Most of the reactions critical to a living cell happen too slowly at normal temperatures to be of any use to the cell. Without enzymes to speed up these reactions, life could not persist. Enzymes do this by binding to the reactant molecules and holding them in such a way as to make the chemical bond-breaking and -forming processes take place more easily. It is important to remember that enzymes do not change whether a reaction is exergonic (spontaneous) or endergonic. This is because they do not change the free energy of the reactants or products. They only reduce the activation energy required for the reaction to go forward (Figure 6). In addition, an enzyme itself is unchanged by the reaction it catalyzes. Once one reaction has been catalyzed, the enzyme is able to participate in other reactions.

The chemical reactants to which an enzyme binds are called the enzyme’s substrates. There may be one or more substrates, depending on the particular chemical reaction. In some reactions, a single reactant substrate is broken down into multiple products. In others, two substrates may come together to create one larger molecule. Two reactants might also enter a reaction and both become modified, but they leave the reaction as two products. The location within the enzyme where the substrate binds is called the enzyme’s active site. The active site is where the “action” happens. Since enzymes are proteins, there is a unique combination of amino acid side chains within the active site. Each side chain is characterized by different properties. They can be large or small, weakly acidic or basic, hydrophilic or hydrophobic, positively or negatively charged, or neutral. The unique combination of side chains creates a very specific chemical environment within the active site. This specific environment is suited to bind to one specific chemical substrate (or substrates).

Active sites are subject to influences of the local environment. Increasing the environmental temperature generally increases reaction rates, enzyme-catalyzed or otherwise. However, temperatures outside of an optimal range reduce the rate at which an enzyme catalyzes a reaction. Hot temperatures will eventually cause enzymes to denature, an irreversible change in the three-dimensional shape and therefore the function of the enzyme. Enzymes are also suited to function best within a certain pH and salt concentration range, and, as with temperature, extreme pH, and salt concentrations can cause enzymes to denature.

For many years, scientists thought that enzyme-substrate binding took place in a simple “lock and key” fashion. This model asserted that the enzyme and substrate fit together perfectly in one instantaneous step. However, current research supports a model called induced fit (Figure 7). The induced-fit model expands on the lock-and-key model by describing a more dynamic binding between enzyme and substrate. As the enzyme and substrate come together, their interaction causes a mild shift in the enzyme’s structure that forms an ideal binding arrangement between enzyme and substrate.

Concept in Action

When an enzyme binds its substrate, an enzyme-substrate complex is formed. This complex lowers the activation energy of the reaction and promotes its rapid progression in one of multiple possible ways. On a basic level, enzymes promote chemical reactions that involve more than one substrate by bringing the substrates together in an optimal orientation for reaction. Another way in which enzymes promote the reaction of their substrates is by creating an optimal environment within the active site for the reaction to occur. The chemical properties that emerge from the particular arrangement of amino acid R groups within an active site create the perfect environment for an enzyme’s specific substrates to react.

The enzyme-substrate complex can also lower activation energy by compromising the bond structure so that it is easier to break. Finally, enzymes can also lower activation energies by taking part in the chemical reaction itself. In these cases, it is important to remember that the enzyme will always return to its original state by the completion of the reaction. One of the hallmark properties of enzymes is that they remain ultimately unchanged by the reactions they catalyze. After an enzyme has catalyzed a reaction, it releases its product(s) and can catalyze a new reaction.

Figure 7. The induced-fit model is an adjustment to the lock-and-key model and explains how enzymes and substrates undergo dynamic modifications during the transition state to increase the affinity of the substrate for the active site.

It would seem ideal to have a scenario in which all of an organism’s enzymes existed in abundant supply and functioned optimally under all cellular conditions, in all cells, at all times. However, a variety of mechanisms ensures that this does not happen. Cellular needs and conditions constantly vary from cell to cell, and change within individual cells over time. The required enzymes of stomach cells differ from those of fat storage cells, skin cells, blood cells, and nerve cells. Furthermore, a digestive organ cell works much harder to process and break down nutrients during the time that closely follows a meal compared with many hours after a meal. As these cellular demands and conditions vary, so must the amounts and functionality of different enzymes.

Since the rates of biochemical reactions are controlled by activation energy, and enzymes lower and determine activation energies for chemical reactions, the relative amounts and functioning of the variety of enzymes within a cell ultimately determine which reactions will proceed and at what rates. This determination is tightly controlled in cells. In certain cellular environments, enzyme activity is partly controlled by environmental factors like pH, temperature, salt concentration, and, in some cases, cofactors or coenzymes.

Enzymes can also be regulated in ways that either promote or reduce enzyme activity. There are many kinds of molecules that inhibit or promote enzyme function, and various mechanisms by which they do so. In some cases of enzyme inhibition, an inhibitor molecule is similar enough to a substrate that it can bind to the active site and simply block the substrate from binding. When this happens, the enzyme is inhibited through competitive inhibition, because an inhibitor molecule competes with the substrate for binding to the active site.

On the other hand, in noncompetitive inhibition, an inhibitor molecule binds to the enzyme in a location other than the active site, called an allosteric site, but still manages to block substrate binding to the active site. Some inhibitor molecules bind to enzymes in a location where their binding induces a conformational change that reduces the affinity of the enzyme for its substrate. This type of inhibition is called allosteric inhibition (Figure 8). Most allosterically regulated enzymes are made up of more than one polypeptide, meaning that they have more than one protein subunit. When an allosteric inhibitor binds to a region on an enzyme, all active sites on the protein subunits are changed slightly such that they bind their substrates with less efficiency. There are allosteric activators as well as inhibitors. Allosteric activators bind to locations on an enzyme away from the active site, inducing a conformational change that increases the affinity of the enzyme’s active site(s) for its substrate(s) (Figure 8).

Figure 8. Allosteric inhibition works by indirectly inducing a conformational change to the active site such that the substrate no longer fits. In contrast, in allosteric activation, the activator molecule modifies the shape of the active site to allow a better fit of the substrate.

Careers in Action

Pharmaceutical Drug Developer

Figure 7. Have you ever wondered how pharmaceutical drugs are developed? (credit: Deborah Austin)

Enzymes are key components of metabolic pathways. Understanding how enzymes work and how they can be regulated are key principles behind the development of many of the pharmaceutical drugs on the market today. Biologists working in this field collaborate with other scientists to design drugs.

Consider statins for example—statins is the name given to one class of drugs that can reduce cholesterol levels. These compounds are inhibitors of the enzyme HMG-CoA reductase, which is the enzyme that synthesizes cholesterol from lipids in the body. By inhibiting this enzyme, the level of cholesterol synthesized in the body can be reduced. Similarly, acetaminophen, popularly marketed under the brand name Tylenol, is an inhibitor of the enzyme cyclooxygenase. While it is used to provide relief from fever and inflammation (pain), its mechanism of action is still not completely understood.

How are drugs discovered? One of the biggest challenges in drug discovery is identifying a drug target. A drug target is a molecule that is literally the target of the drug. In the case of statins, HMG-CoA reductase is the drug target. Drug targets are identified through painstaking research in the laboratory. Identifying the target alone is not enough scientists also need to know how the target acts inside the cell and which reactions go awry in the case of disease. Once the target and the pathway are identified, then the actual process of drug design begins. In this stage, chemists and biologists work together to design and synthesize molecules that can block or activate a particular reaction. However, this is only the beginning: If and when a drug prototype is successful in performing its function, then it is subjected to many tests from in vitro experiments to clinical trials before it can get approval from the U.S. Food and Drug Administration to be on the market.

Many enzymes do not work optimally, or even at all, unless bound to other specific non-protein helper molecules. They may bond either temporarily through ionic or hydrogen bonds, or permanently through stronger covalent bonds. Binding to these molecules promotes optimal shape and function of their respective enzymes. Two examples of these types of helper molecules are cofactors and coenzymes. Cofactors are inorganic ions such as ions of iron and magnesium. Coenzymes are organic helper molecules, those with a basic atomic structure made up of carbon and hydrogen. Like enzymes, these molecules participate in reactions without being changed themselves and are ultimately recycled and reused. Vitamins are the source of coenzymes. Some vitamins are the precursors of coenzymes and others act directly as coenzymes. Vitamin C is a direct coenzyme for multiple enzymes that take part in building the important connective tissue, collagen. Therefore, enzyme function is, in part, regulated by the abundance of various cofactors and coenzymes, which may be supplied by an organism’s diet or, in some cases, produced by the organism.

Free Energy Embodies the First and Second Laws

Free energy (G) is the energy available (or required) to do work in a given system. If a given system releases free energy, then it can do work. Conversely, if it absorbs free energy, then work can be done on it.

Let's return to the example of marbles being held in a hand. We will define the system as the person holding the marbles and the marbles themselves. When the marbles are held, they are relatively ordered (they have low entropy) but unstable (simply opening the hand will cause a spontaneous change in the system). The potential energy of the marbles is also relatively high. When the hand is opened, however, several things happen. First, potential energy is converted to kinetic energy. Second, friction is produced and released in the form of heat as the marbles fall through the air. Third, the system becomes disordered as the marbles bounce around (entropy increases). If the person walks around and picks up the marbles, then the thermodynamic state changes again. First, the kinetic energy exerted by the person picking up the marbles is converted to gravitational potential energy as the marbles become elevated above the ground, and second, the marbles become more ordered (less entropic). Importantly, the ordered state can be restored, but energy (in the form of the person picking up the marbles) is required to order matter.

The change in free energy (delta G) is endergonic if energy enters the system, and exergonic if it leaves the system. Moreover, an exergonic reaction is unstable, has a negative delta G, and is therefore a spontaneous reaction.

Lastly, in this example one can see why energy flow is not 100% efficient as the marbles fall through the air there is a production of frictional heat (which, in this example, does no useful work and represents waste). All energy transfers have some inefficiencies, which is why reactions do not transduce 100% of the available energy.

Organisms can only live at the expense of free energy (G). The free energy changes (delta G) associated with life's metabolic energy involve the movement of matter. This free energy comes from a series of metabolic reactions that result in work being done at the molecular level (the movement of electrons, atoms, or molecules). Recall the relationship above, between free energy and stability a given reaction (a system) that has the potential to do a lot of work (release a lot of free energy) is inherently unstable it typically has a low relative entropy and tends to change spontaneously to a more stable, disordered state. In fact, the concept of spontaneity actually defines whether free energy is made available to do work (or if work is required).

Free energy is more than a change in entropic state because each given system has a certain amount of total energy. However, not all of this total energy is available to do work. Free energy is a function of the total energy change of a system and the entropic change.

Review Questions

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, iii. just before it reaches the end of one cycle (just before instant i.).

  1. i. potential and kinetic, ii. potential and kinetic, iii. kinetic
  2. i. potential, ii. potential and kinetic, iii. potential and kinetic
  3. i. potential, ii. kinetic, iii. potential and kinetic
  4. i. potential and kinetic, ii. kinetic iii. kinetic

Which of the following comparisons or contrasts between endergonic and exergonic reactions is false?

  1. Endergonic reactions have a positive ∆G and exergonic reactions have a negative ∆G
  2. Endergonic reactions consume energy and exergonic reactions release energy
  3. Both endergonic and exergonic reactions require a small amount of energy to overcome an activation barrier
  4. 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?

  1. Compare the ∆G values between the two reactions
  2. Compare their reaction rates
  3. Compare their ideal environmental conditions
  4. Compare the spontaneity between the two reactions


In mechanistic terms, an enzyme-catalyzed reaction is a multistep reaction. When the energy profile of such a reaction is depicted, it is convenient to explicitly state the prevailing conditions (state of the system). If these conditions allow the overall reaction to proceed in the forward direction, then it must be noted that every single step of the reaction mechanism is a spontaneous process, and therefore it must exhibit a negative free energy change. These considerations must be properly reflected in the progression profile of the reaction.

Gibbs free energy reaction coordinate profiles found in some textbooks. The energy diagram for a reaction model consisting of one enzyme, one substrate, and one product is depicted in many books where it is compared with that for the uncatalyzed reaction. The survey of several Biochemistry textbooks reveals a high diversity of profiles for the same process. A, C, and D are adapted from Refs. 5 , 3 , and 2 , respectively. B is adapted from Ref. 6 or 7 . Substrate, intermediates (ES and EP), and product are traced in black, while the transition states are depicted in blue and red for the uncatalyzed and catalyzed reaction mechanism, respectively. Symbols are explained in the text.

Thermodynamic pits and hills. This figure is often used to illustrate the contribution to the catalysis of a weak binding of the enzyme to its substrate [ 5 , 8 , 10 and 11 ]. A, Low values for the Michaelis constant may make difficult the catalysis by decreasing ΔGgs and thus increasing ΔGc ≠ (see text for details). However, high Km values cannot drive reactions against thermodynamic potential as it is implicitly suggested by B.

Weak binding of substrates to enzymes. The intrinsic binding energy of the ES complex is compensated to some extent by entropy loss due to the binding of E and S. If the enzyme bound the substrate too tightly, the activation barrier would be comparable to the activation barrier of the non-enzymatic reaction (A). When the enzyme meets the Michaelis-Menten condition for rapid equilibrium (k−1k2), then ΔGgs tends toward zero, which is kinetically optimal (B).