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2.15: Introduction to Atomic Bonds - Biology


What you’ll learn to do: Classify different types of atomic bonds

When atoms bond together, they create molecules: a sodium atom bonds with a chlorine atom to create salt (sodium chloride), two hydrogen atoms bond with an oxygen atom to create water (hydrogen dioxide). However, not all atomic bonds are the same; in fact salt and water are created with two very different types of bonds (ionic and polar covalent bonds respectively).

The different types of bonds (ionic, polar covalent, and non-polar covalent bonds) behave differently, and these differences have an impact on the molecules they create. Understanding the types of bonds that create living things can help us understand those living things themselves.


Atomic bonds

Once the way atoms are put together is understood, the question of how they interact with each other can be addressed—in particular, how they form bonds to create molecules and macroscopic materials. There are three basic ways that the outer electrons of atoms can form bonds:

The first way gives rise to what is called an ionic bond. Consider as an example an atom of sodium, which has one electron in its outermost orbit, coming near an atom of chlorine, which has seven. Because it takes eight electrons to fill the outermost shell of these atoms, the chlorine atom can be thought of as missing one electron. The sodium atom donates its single valence electron to fill the hole in the chlorine shell, forming a sodium chloride system at a lower total energy level.

An atom that has more or fewer electrons in orbit than protons in its nucleus is called an ion. Once the electron from its valence shell has been transferred, the sodium atom will be missing an electron it therefore will have a positive charge and become a sodium ion. Simultaneously, the chlorine atom, having gained an extra electron, will take on a negative charge and become a chlorine ion. The electrical force between these two oppositely charged ions is attractive and locks them together. The resulting sodium chloride compound is a cubic crystal, commonly known as ordinary table salt.

The second bonding strategy listed above is described by quantum mechanics. When two atoms come near each other, they can share a pair of outermost electrons (think of the atoms as tossing the electrons back and forth between them) to form a covalent bond. Covalent bonds are particularly common in organic materials, where molecules often contain long chains of carbon atoms (which have four electrons in their valence shells).

Finally, in some materials each atom gives up an outer electron that then floats freely—in essence, the electron is shared by all of the atoms within the material. The electrons form a kind of sea in which the positive ions float like marbles in molasses. This is called the metallic bond and, as the name implies, it is what holds metals together.

There are also ways for atoms and molecules to bond without actually exchanging or sharing electrons. In many molecules the internal forces are such that the electrons tend to cluster at one end of the molecule, leaving the other end with a positive charge. Overall, the molecule has no net electric charge—it is just that the positive and negative charges are found at different places. For example, in water (H2O) the electrons tend to spend most of their time near the oxygen atom, leaving the region of the hydrogen atoms with a positive charge. Molecules whose charges are arranged in this way are called polar molecules. An atom or ion approaching a polar molecule from its negative side, for example, will experience a stronger negative electric force than the more-distant positive electric force. This is why many substances dissolve in water: the polar water molecule can pull ions out of materials by exerting electric forces. A special case of polar forces occurs in what is called the hydrogen bond. In many situations, when hydrogen forms a covalent bond with another atom, electrons move toward that atom, and the hydrogen acquires a slight positive charge. The hydrogen, in turn, attracts another atom, thereby forming a kind of bridge between the two. Many important molecules, including DNA, depend on hydrogen bonds for their structure.

Finally, there is a way for a weak bond to form between two electrically neutral atoms. Dutch physicist Johannes van der Waals first theorized a mechanism for such a bond in 1873, and it is now known as van der Waals forces. When two atoms approach each other, their electron clouds exert repulsive forces on each other, so that the atoms become polarized. In such situations, it is possible that the electrical attraction between the nucleus of one atom and the electrons of the other will overcome the repulsive forces between the electrons, and a weak bond will form. One example of this force can be seen in ordinary graphite pencil lead. In this material, carbon atoms are held together in sheets by strong covalent bonds, but the sheets are held together only by van der Waals forces. When a pencil is drawn across paper, the van der Waals forces break, and sheets of carbon slough off. This is what creates the dark pencil streak.


Chemical bonds and their properties

The concept of the chemical bond lies at the very core of Chemistry it is what enables about one hundred elements to form the more than fifty million known chemical substances that make up our physical world. Before we get into the theory of chemical bonding, we need to define what we are talking about:
Exactly what is a chemical bond?
And what observable properties can we use to distingish one kind of bond from another? This is the first of ten lessons that will help familiarize you with the fundamental concepts of this very broad subject.

You probably learned some time ago that chemical bonds are what hold atoms together to form the more complicated aggregates that we know as molecules and extended solids. Chemists talk about bonds all the time, and draw pictures of them as lines joining atom symbols. Teachers often identify them as the little sticks that connect the spheres that represent atoms in a plastic molecular model. So it's not surprising that we sometimes tend to think of chemical bonds as “things”. But no one has ever seen a chemical bond, and there is no reason to believe that they really even exist as physical objects.

"SOMETIMES IT SEEMS to me that a bond between two atoms has become so real, so tangible, so friendly, that I can almost see it. Then I awake with a little shock, for a chemical bond is not a real thing. It does not exist. No one has ever seen one. No one ever can. It is a figment of our own imagination."

C.A. Coulson (1910-1974) was an English theoretical chemist who played a central role in the development of quantum theories of chemical bonding.

It is probably more useful to regard a chemical bond as an effect that causes certain atoms to join together to form enduring structures that have unique physical and chemical properties.

So although the "chemical bond" (as a physical object) may be no more than a convenient fiction, chemical bonding, which leads to the near-infinity of substances (31 million in mid-2007), lies at the very core of chemistry.

The forces that hold bonded atoms together are basically just the same kinds of electrostatic attractions that bind the electrons of an atom to its positively-charged nucleus

This is the most important fact about chemical bonding that you should know, but it is not of itself a workable of bonding because it does not describe the conditions under which bonding occurs, nor does it make useful predictions about the properties of the bonded atoms.

Our views of what constitutes chemical bonding are still evolving, according to a 2007 article in Chemical and Engineering News(85 37-40). This "buckyball-and-mitt" synthesized in 2007 by Andrzej Sygula is a case in point. The buckyball C60 resides in the C60H28"buckybowl". There are no traditional "chemical bonds" between the ball and the mit!

image from C&EN 85 (13) 2008

Most people think of molecules as the particles that result when atoms become joined together in some way. This conveys the general picture, but a somewhat better definition that we will use in these lessons is

A more restrictive definition distinguishes between a "true" molecule that exists as an independent particle, and an that can only be represented by its simplest formula. Methane, CH4, is an example of the former, while sodium chloride, which does not contain any discrete NaCl units, is the most widely-known extended solid. But because we want to look at chemical bonding in the most general way, we will avoid making this distinction here except in a few special cases. In order to emphasize this "aggregate of atoms" definition, we will often use terms such as "chemical species" and "structures" in place of "molecules" in this lesson.

The definition written above is an operational one that is, it depends on our ability to observe and measure the molecule's properties. Clearly, this means that the molecule must retain its identity for a period of time long enough to carry out these observations. For most of the molecules of chemical interest, this presents no difficulty. But it does happen that some structures that we can write formulas for, such as He2, have such brief lives that no significant properties have been observed. So to some extent, what we consider to be a molecule depends on the technology we use to observe them, and this will necessarily change with time.

Structure, structure, structure!

And what are those properties that characterize a particular kind of molecule and distinguish it from others? Just as real estate is valued by "location, location, location", the identity of a chemical species is defined by its . In its most fundamental sense, the structure of a molecule is specified by the identity of its constituent atoms and the sequence in which they are joined together, that is, by the . This, in turn, defines the &mdash the spatial relationship between the bonded atoms.

The importance of bonding connectivity is nicely illustrated by the structures of the two compounds ethanol and dimethyl ether, both of which have the simplest formula C2H6O. The structural formulas reveal the very different connectivities of these two molecules whose physical and chemical properties are quite different:

The precise definition of bonding energy is described in another lesson and is not important here. For the moment you only need to know that in any stable structure, the potential energy of its atoms is lower than that of the individual isolated atoms. Thus the formation of methane from its gaseous atoms (a reaction that cannot be observed under ordinary conditions but for which the energetics are known from indirect evidence)

is accompanied by the release of heat, and is thus an exothermic process. The quantity of heat released is related to the stability of the molecule. The smaller the amount of energy released, the more easily can the molecule absorb thermal energy from the environment, driving the above reaction in reverse and leading to the molecule's decomposition. A highly stable molecule such as methane must be subjected to temperatures of more than 1000°C for significant decomposition to occur. But the noble-gas molecule KrF2 is so weakly bound that it decomposes even at 0°C, and the structure He2 has never been observed. If a particular arrangement of atoms is too unstable to reveal its properties at any achievable temperature, then it does not qualify to be called a molecule.

There are many molecules that are energetically stable enough to meet the above criterion, but are so that their lifetimes are too brief to make their observation possible. The molecule CH3, methyl, is a good example: it can be formed by electrical discharge in gaseous CH4, but it is so reactive that it combines with almost any molecule it strikes (even another CH3) within a few collisions. It was not until the development of spectroscopic methods (in which a molecule is characterized by the wavelengths of light that it absorbs or emits) that methyl was recognized as a stable albeit shamelessly promiscuous molecule that is an important intermediate in many chemical processes ranging from flames to atmospheric chemistry.

Chemical species are traditionally represented by such as the ones for phosphoric acid, H3PO4, which we show here. The lines, of course, represent the "chemical bonds" of the molecule.
More importantly, the structural formula of a molecule defines its , as was illustrated in the comparison of ethanol and dimethyl ether shown in the previous section.

One limitation of such formulas is that they are drawn on a two-dimensional surface, whereas most molecules have a three-dimensional shape.

The wedge-shaped lines in the structural formula on the right are one way of indicating which bonds extend above or below the viewing plane, providing a kind of pseudo-3D view. You will probably be spared having to learn this convention until you get into Organic Chemistry.

Three-dimensional models (either real plastic ones or images that incorporate perspective and shading) reveal much more about a molecule's structure. The ball-and-stick and space-filling renditions are widely employed, but each has its limitations, as seen in the following examples that compare different ways of depicting the structures of the same two molecules:

Simple structural formulas in two dimensions show the molecule's connectivity, but nothing more.

&uarr This simple structural formula of methane, CH4, projects its 3-dimentional structure onto a 2D surface.

&uarr The structural formula of ascorbic acid (vitamin C) is commonly enhanced with wedge-shaped bonds to show that they extend above the plane of the paper or screen.

Ball-and-stick models show 3-dimensional views of the "chemical bonds" and their geometry. although with the individual atoms unrealisticly separated.

This image correctly expresses the tetrahedral coordination of the four C—H bonds.

Space-filling models don't attempt to depict the bonds, but show the relative sizes of the atoms and general shape of the molecule, at the expense of hiding some of the atoms.


Note how this shows CH4 to be roughly spherical.

Finally, we get to see one! In 2009, IBM scientists in Switzerland succeeded in imaging a real molecule, using a technique known as atomic force microscopy in which an atoms-thin metallic probe is drawn ever-so-slightly above the surface of an immobilized pentacene molecule cooled to nearly absolute zero. In order to improve the image quality, a molecule of carbon monoxide was placed on the end of the probe.

The image produced by the AFM probe is shown at the very bottom. What is actually being imaged is the surface of the electron clouds of the molecule, which consists of five fused hexagonal rings of carbon atoms with hydrogens on its periphery. The tiny bumps that correspond to these hydrogen atoms attest to the remarkable resolution of this experiment.

The original article was published in Science magazine see here for an understandable account of this historic work.

The purpose of rendering a molecular structure in a particular way is not to achieve "realism" (whatever that might be), but rather to convey useful information of some kind. Modern computer rendering software takes its basic data from various kinds of standard structural databases which are compiled either from experimental X-ray scattering data, or are calculated from theory.


[image]

Example: Caffeine

Caffeine's molecular formula, C8H10N4O2, tells us its composition, but conveys no information about how how its 24 atoms are connected, and thus how it differs from any number of other compounds having the same formula. In order to reveal its connectivity, we use its .

Although the structure at the right does not look like
the one on the cup, their connectivities are identical,
so the two structures are equivalent. As an exercise,
see if you are able to confirm this.

For most "ordinary" chemistry, the structural formula is all we need. Suppose, however, that you wish to understand more about how caffeine exerts its stimulating effect on the body. Like many drugs, caffeine binds to specific sites on proteins the efficacy of this binding is usually highly dependent on the both the shape of the drug and the manner in which electric charge is distributed over this shape.

In this context, it is important to understand that the outer "surface" of a molecule is defined by the veil of negative charge that originates in the valence electrons of the atoms but which tends to be spread over the entire molecule to a distance that can significantly affect with neighboring molecules.

For this purpose, one employs a mixture of chemical intuition and molecular modeling computer software to generate an image such as the one shown here.

In this depiction of Caffeine, the atoms are coded by color: white=H, red=O, light blue=C, dark blue=N.

Example: Sucrose

Sucrose &mdash ordinary "sugar" — occurs naturally in fruits and vegetables, as well as in everyone's kitchen and in all too many foods and beverages. Its structural formula shows that it is really a "double sugar" (a disaccharide) in which two monosaccharides, glucose and fructose, are joined together.

The overall shape and distribution of electric charge at the surface enable it to bind to the sweetness receptors on the tongue, and — more importantly to the enzyme that catalyzes the hydrolysis reaction that breaks the sucrose into its two monosaccharide components, releasing the glucose which fuels our body's cells.

Knowing the properties of molecular surfaces is vitally important to understanding any process that depends on one molecule remaining in physical contact with another. Catalysis is one example, but one of the main interests at the present time is biological signalling, in which a relatively small molecule binds to or "docks" with a receptor site on a much larger one, often a protein. Sophisticated molecular modeling software such as was used to produce these images is now a major tool in many areas of research biology.
[Images: left, right ]

Visualizing very large molecules such as carbohydrates and proteins that may contain tens of thousands of atoms presents obvious problems. The usual technique is to simplify parts of the molecule, representing major kinds of extended structural units by shapes such as ribbons or tubes which are twisted or bent to approximate their conformations. These are then gathered to reveal the geometrical relations of the various units within the overall structure. Individual atoms, if shown at all, are restricted to those of special interest.


Netropsin is a naturally-occurring antibiotic that binds to specific sites on DNA. The netropsin itself is rendered as a ball-and-stick model with an extended molecular surface. The tag-like "attachments" on the more colorful double-stranded DNA chain represent sugar molecules and nucleotide bases. [source]

Some of these images can stand as artistic creations in their own right. This seems to be especially true of those that render special aspects of molecular surfaces the two on the left below could well be mistaken for paintings by Jean Miró and Salvador Dalí, respectively.


This creature making its graceful descent illustrates some of the topological forms involved in molecular surface problems. This image and the one at the left are reproduced here with the kind permission of Prof. Sanner.


Like an insect caught in a spider's web, this alien invader named Heme is fighting to extricate itself from the Globin in which he is entrapped.

Study of the surface properties of large molecules is crucial for understanding how proteins, carbohydrates, and DNA interact with smaller molcules, especially those involved in transport of ions and small molecule across cell membranes, immune-system behavior, and signal transduction processes such as the "turning on" of genes.

See here for links to a wide variety of sources relating to visualization and molecular modeling.

When we talk about the properties of a particular chemical bond, we are really discussing the relationship between two adjacent atoms that are part of the molecule. Diatomic molecules are of course the easiest to study, and the information we derive from them helps us interpret various kinds of experiments we carry out on more complicated molecules.

It is important to bear in mind that the exact properties of a specific kind of bond will be determined in part by the nature of the other bonds in the molecule thus the energy and length of the C–H bond will be somewhat dependent on what other atoms are connected to the carbon atom. Similarly, the C-H bond length can vary by as much a 4 percent between different molecules. For this reason, the values listed in tables of bond energy and bond length are usually averages taken over a variety of compounds that contain a specific atom pair..

In some cases, such as C—O and C—C, the variations can be much greater, approaching 20 percent. In these cases, the values fall into groups which we interpret as representative of single- and multiple bonds: double, and triple.

The energy of a system of two atoms depends on the distance between them. At large distances the energy is zero, meaning “no interaction”. At distances of several atomic diameters attractive forces dominate, whereas at very close approaches the force is repulsive, causing the energy to rise. The attractive and repulsive effects are balanced at the minimum point in the curve. Plots that illustrate this relationship are known as , and they are quite useful in defining certain properties of a chemical bond.

The internuclear distance at which the potential energy minimum occurs defines the bond length. This is more correctly known as the equilibrium bond length, because thermal motion causes the two atoms to vibrate about this distance. In general, the stronger the bond, the smaller will be the bond length.

Attractive forces operate between all atoms, but unless the potential energy minimum is at least of the order of RT, the two atoms will not be able to withstand the disruptive influence of thermal energy long enough to result in an identifiable molecule. Thus we can say that a chemical bond exists between the two atoms in H2. The weak attraction between argon atoms does not allow Ar2 to exist as a molecule, but it does give rise to the that holds argon atoms together in its liquid and solid forms.

Potential energy and kinetic energy Quantum theory tells us that an electron in an atom possesses K as well as P, so the total energy E is always the sum of the two: E = P + K. The relation between them is surprisingly simple: K = ל.5 P. This means that when a chemical bond forms (an exothermic process with &DeltaE < 0), the decrease in potential energy is accompanied by an increase in the kinetic energy (embodied in the momentum of the bonding electrons), but the magnitude of the latter change is only half as much, so the change in potential energy always dominates. The –&DeltaE has half the magnitude of the fall in potential energy.

How bond energies are measured

Bond energies are usually determined indirectly from thermodynamic data, but there are two main experimental ways of measuring them directly:

1. The direct thermochemical method involves separating the two atoms by an electrical discharge or some other means, and then measuring the heat given off when they recombine. Thus the energy of the C—C single bond can be estimated from the heat of the recombination reaction between methyl radicals, yielding ethane:

Although this method is simple in principle, it is not easy to carry out experimentally. The highly reactive components must be prepared in high purity and in a stream of moving gas.

2. The spectroscopic method is based on the principle that absorption of light whose wavelength corresponds to the bond energy will often lead to the breaking of the bond and dissociation of the molecule. For some bonds, this light falls into the green and blue regions of the spectrum, but for most bonds ultraviolet light is required. The experiment is carried out by observing the absorption of light by the substance being studied as the wavelength is decreased. When the wavelength is sufficiently small to break the bond, a characteristic change in the absorption pattern is observed.

Spectroscopy is quite easily carried out and can yield highly precise results, but this method is only applicable to a relatively small number of simple molecules. The major problem is that the light must first be absorbed by the molecule, and relatively few molecules happen to absorb light of a wavelength that corresponds to a bond energy.

Experiments carried out on diatomic molecules such as O2 and CS yield unambiguous values of bond energy, but for more complex molecules there are complications. For example, the heat given off in the CH3 combination reaction written above will also include a small component that represents the differences in the energies of the C-H bonds in methyl and in ethane. These can be corrected for by experimental data on reactions such as

By assembling a large amount of experimental information of this kind, a consistent set of average bond energies can be obtained (see table below.) The energies of double bonds are greater than those of single bonds, and those of triple bonds are higher still.

One can often get a very good idea of how much heat will be absorbed or given off in a reaction by simply finding the difference in the total bond energies contained in the reactants and products. The strength of an individual bond such as O–H depends to some extent on its environment in a molecule (that is, in this example, on what other atom is connected to the oxygen atom), but tables of "average" energies of the various common bond types are widely available and can provide useful estimates of the quantity of heat absorbed or released in many chemical reaction.

The bond energy units in the above table are kilojoules per mole.

As an example, consider the reaction of chlorine with methane to produce dichloromethane and hydrogen chloride:

In this reaction, two C–Cl bonds and two H–Cl bonds are broken, and two new C–Cl and H–Cl bonds are formed. The net change associated with the reaction is

which comes to 𤫀 kJ per mole of methane this agrees quite well with the observed heat of reaction, which is 𤪺 kJ/mol.

The length of a chemical bond the distance between the centers of the two bonded atoms (the .) Bond lengths have traditionally been expressed in Ångstrom units, but picometers are now preferred (1Å = 10 -8 cm = 100 pm.) Bond lengths are typically in the range 1-2 Å or 100-200 pm. Even though the bond is vibrating, equilibrium bond lengths can be determined experimentally to within ±1 pm.

Bond lengths depend mainly on the sizes of the atoms, and secondarily on the bond strengths, the stronger bonds tending to be shorter. Bonds involving hydrogen can be quite short The shortest bond of all, H–H, is only 74 pm. Multiply-bonded atoms are closer together than singly-bonded ones this is a major criterion for experimentally determining the multiplicity of a bond. This trend is clearly evident in the above plot which depicts the sequence of carbon-carbon single, double, and triple bonds.

The most common method of measuring bond lengths in solids is by analysis of the diffraction or scattering of X-rays when they pass through the regularly-spaced atoms in the crystal. For gaseous molecules, neutron- or electron-diffraction can also be used.

The complete structure of a molecule requires a specification of the coordinates of each of its atoms in three-dimensional space. This data can then be used by computer programs to construct visualizations of the molecule as discussed above. One such visualization of the water molecule, with bond distances and the HOH bond angle superimposed on a space-filling model, is shown here. (It is taken from an excellent reference source on water). The colors show the results of calculations that depict the way in which electron charge is distributed around the three nuclei.


All molecules are in a constant state of internal motion this animation shows some vibrational modes of benzene, C6H6. [source]

When an atom is displaced from its equilibrium position in a molecule, it is subject to a restoring force which increases with the displacement. A spring follows the same law (Hooke’s law) a chemical bond is therefore formally similar to a spring that has weights (atoms) attached to its two ends. A mechanical system of this kind possesses a natural vibrational frequency which depends on the masses of the weights and the stiffness of the spring. These vibrations are initiated by the thermal energy of the surroundings chemically-bonded atoms are never at rest at temperatures above absolute zero.

On the atomic scale in which all motions are quantized, a vibrating system can possess a series of vibrational frequencies, or states. These are depicted by the horizontal lines in the potential energy curve shown here. Notice that the very bottom of the curve does not correspond to an allowed state because at this point the positions of the atoms are precisely specified, which would violate the uncertainty principle. The lowest-allowed, or ground vibrational state is the one denoted by 0, and it is normally the only state that is significantly populated in most molecules at room temperature. In order to jump to a higher state, the molecule must absorb a photon whose energy is equal to the distance between the two states.

For ordinary chemical bonds, the energy differences beween these natural vibrational frequencies correspond to those of . Each wavelength of infrared light that excites the vibrational motion of a particular bond will be absorbed by the molecule. In general, the stronger the bond and the lighter the atoms it connects, the higher will be its natural stretching frequency and the shorter the wavelength of light absorbed by it. Studies on a wide variety of molecules have made it possible to determine the wavelengths absorbed by each kind of bond (See here for a brief list.) By plotting the degree of absorption as a function of wavelength, one obtains the of the molecule which allows one to "see" what kinds of bonds are present.

Infrared spectrum of alcohol The low points in the plot below indicate the frequencies of infrared light that are absorbed by ethanol (ethyl alcohol),
CH3CH2OH. Notice how stretching frequencies involving hydrogen are higher, reflecting the smaller mass of that atom. Only the most prominent absorption bands are noted here.

Now that you know something about bond stretching vibrations,
you can impress your friends by telling them why water is blue

Actual infrared spectra are complicated by the presence of more complex motions (stretches involving more than two atoms, wagging, etc.), and absorption to higher quantum states (overtones), so infrared spectra can become quite complex. This is not necessarily a disadvantage, however, because such spectra can serve as a "fingerprint" that is unique to a particular molecule and can be helpful in identifying it. Largely for this reason, infrared spectrometers are standard equipment in most chemistry laboratories.

The aspect of bond stretching and bending frequencies that impacts our lives most directly is the way that some of the gases of the atmosphere absorb infrared light and thus affect the heat balance of the Earth. Owing to their symmetrical shapes, the principal atmospheric components N2 and O2 do not absorb infrared light, but the minor components water vapor and carbon dioxide are strong absorbers, especially in the long-wavelength region of the infrared. Absorption of infrared light by a gas causes its temperature to rise, so any source of infrared light will tend to warm the atmosphere this phenomenon is known as the .

The incoming radiation from the Sun (which contains relatively little long-wave infrared light) passes freely through the atmosphere and is absorbed by the Earth's surface, warming it up and causing it to re-emit some of this energy as long-wavelength infrared. Most of the latter is absorbed by the H2O and CO2 , the major greenhouse gasis in the unpolluted atmosphere, effectively trapping the radiation as heat. Thus the atmosphere is heated by the Earth, rather than by direct sunlight. Without the “” in the atmosphere, the Earth's heat would be radiated away into space, and our planet would be too cold for life.

Radiation balance of the Earth In order to maintain a constant average temperature, the quantity of radiation (sunlight) absorbed by the surface must be exactly balanced by the quantity of long-wavelength infrared emitted by the surface and atmosphere and radiated back into space. Atmospheric gases that absorb this infrared light (depicted in red on the right part of this diagram) partially block this emission and become warmer, raising the Earth's temperature. This diagram is from the U. of Oregon Web page referenced below.

Since the beginning of the Industrial Revolution in the 19th century, huge quantities of additional greenhouse gases have been accumulating in the atmosphere. Carbon dioxide from fossil fuel combustion has been the principal source, but intensive agriculture also contributes significant quantities of methane (CH4) and nitrous oxide (N2O) which are also efficient far-infrared absorbers. The measurable increase in these gases is believed by many to be responsible for the increase in the average temperature of the Earth that has been noted over the past 50 years— a trend that could initiate widespread flooding and other disasters if it continues.

Make sure you thoroughly understand the following essential concepts that have been presented above.

  • How would you define a ?
  • What is meant by the of a molecule? What additional information might be needed in order to specify its structure?
  • Explain the difference between and , and how these factors might prevent a given structure from existing long enough to qualify as a .
  • Sketch out a potential energy curve for a typical diatomic molecule, and show how it illustrates the and .
  • Explain how the heat released or absorbed in a chemical reaction can be related to the bond energies of the reactants and products.
  • State the major factors that determine the distance between two bonded atoms.
  • Describe, in a general way, how the of a substance can reveal details about its molecular structure.

© 2004-2016 by Stephen Lower - last modified 2017-10-27

For information about this Web site or to contact the author,
please see the Chem1 Virtual Textbook home page.

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Chem1 Chemical fonding - part 1 is an overview of Chemical bonds and their properties and serves as an introduction to the subject of chemical bonding at a level suitable for a course in General Chemistry . It is part of the General Chemistry Virtual Textbook , a free, online reference textbook for General Chemistry by Stephen Lower of

This chapter covers the following topics: Definition of a chemical bond and of a molecule, structural formulas, visualization of complex structures, potential energy curves, bond energies, bond lengths, infrared absorption spectra, greenhouse gas warmng. It can be accessed directly at http://www.chem1.com/acad/webtext/chembond/cb01.html .

This material is directed mainly at the first-year college level, but much of it is also suitable for high-school students. It is licensed under a Creative Commons Attribution 3.0 Unported License .


Covalent Bonds

Another type of strong chemical bond between two or more atoms is a covalent bond. These bonds form when an electron is shared between two elements and are the strongest and most common form of chemical bond in living organisms. Covalent bonds form between the elements that make up the biological molecules in our cells. Unlike ionic bonds, covalent bonds do not dissociate in water.

Interestingly, chemists and biologists measure bond strength in different ways. Chemists measure the absolute strength of a bond (the theoretical strength) while biologists are more interested in how the bond behaves in a biological system, which is usually aqueous (water-based). In water, ionic bonds come apart much more readily than covalent bonds, so biologists would say that they are weaker than covalent bonds. If you look in a chemistry textbook, you’ll see something different. This is a great example of how the same information can lead to different answers depending on the perspective that you’re viewing it from.

The hydrogen and oxygen atoms that combine to form water molecules are bound together by covalent bonds. The electron from the hydrogen atom divides its time between the outer shell of the hydrogen atom and the incomplete outer shell of the oxygen atom. To completely fill the outer shell of an oxygen atom, two electrons from two hydrogen atoms are needed, hence the subscript “2” in H 2 O. The electrons are shared between the atoms, dividing their time between them to “fill” the outer shell of each. This sharing is a lower energy state for all of the atoms involved than if they existed without their outer shells filled.

There are two types of covalent bonds: polar and nonpolar. Nonpolar covalent bonds form between two atoms of the same element or between different elements that share the electrons equally. For example, an oxygen atom can bond with another oxygen atom to fill their outer shells. This association is nonpolar because the electrons will be equally distributed between each oxygen atom. Two covalent bonds form between the two oxygen atoms because oxygen requires two shared electrons to fill its outermost shell. Nitrogen atoms will form three covalent bonds (also called triple covalent) between two atoms of nitrogen because each nitrogen atom needs three electrons to fill its outermost shell. Another example of a nonpolar covalent bond is found in the methane (CH 4 ) molecule. The carbon atom has four electrons in its outermost shell and needs four more to fill it. It gets these four from four hydrogen atoms, each atom providing one. These elements all share the electrons equally, creating four nonpolar covalent bonds (Figure 3).

In a polar covalent bond , the electrons shared by the atoms spend more time closer to one nucleus than to the other nucleus. Because of the unequal distribution of electrons between the different nuclei, a slightly positive (δ+) or slightly negative (δ–) charge develops. The covalent bonds between hydrogen and oxygen atoms in water are polar covalent bonds. The shared electrons spend more time near the oxygen nucleus, giving it a small negative charge, than they spend near the hydrogen nuclei, giving these molecules a small positive charge.

Figure 3 The water molecule (left) depicts a polar bond with a slightly positive charge on the hydrogen atoms and a slightly negative charge on the oxygen. Examples of nonpolar bonds include methane (middle) and oxygen (right).


Constructing a Bohr Model of an Atom

Understanding electron configuration is key to understanding chemical reactivity. While we can determine the number of electrons from the atomic number, their configuration needs some additional explanation. Electrons tend radiate around the nucleus at distinct distances, known as orbitals. The structure of the periodic table can help you determine the number of orbitals an element has and how many electrons are in each orbital.

The periodic table is broken into periods (rows) and groups (columns). Each period represents the number of electron orbitals an atom of an element has. We can also determine how those electrons are arranged based on the element’s position on the periodic table. Electrons fill the inner orbitals before adding new orbitals. For example, Lithium (Li) has two orbitals because it is in the second period (second row). Since Lithium’s atomic number is 3, we know that it has 3 electrons. The first two electrons fill the first orbital, and the last orbital (known as the valence orbital) has one electron (known as the valence electron). In fact all elements that are in the first column (Group 1) have one valence electron, but different numbers of orbitals. This gives them similar chemical properties.

Figure 2. Various ways of representing an atom of the element, lithium. a) Lithium as represented on the periodic table. Lithium has an atomic number of 3 (indicating it contains 3 protons and three electrons) and an atomic mass of 6.94. Rounding the atomic mass to 7, you can determine that lithium has 4 neutrons. b) Lithium as represented by the Bohr model of an atom. The protons and neutrons are represented in the middle of the atom, or nucleus. Unlike helium, lithium has two electron orbitals. You know this because it is in the second row, or period. The inner orbitals must fill with electrons before outer orbitals acquire electrons. The number of possible electrons per orbital are equal to the number of elements within a period. The first period have two elements, and therefore only two electrons can occupy the first orbital. Lithium has three electrons. Therefore the first two electrons fill the first orbital and the remaining electron occupies the second orbital. The outside orbital is known as the valence orbital, and the electrons of the valence orbital are known as valence electrons. There are 8 elements in the second period indicating that the second orbital can hold up to 8 electrons. c) The electron dot diagram of lithium. Lithium has one valence electron, which is represented by a single dot.

Fig. 3. The Periodic Table of Elements. For this lab, we will only concentrate on the first three rows (or periods). Each period represents the number of electron orbitals an atom of an element has. For example, element 6, carbon (C) is in the second period indicating it has two orbitals. We can also determine how those electrons are arranged based on the element’s position on the periodic table. The number of electrons per orbital is equal to the number of elements within a period. The first period has two elements, indicating the first orbital can have up to two electrons. The second orbital can have up to eight. Electrons fill the inner orbitals before adding new orbitals. To determine the number of valence electrons (electrons in the outer orbital), you can simple count from left to right on the period the element is in. For example, carbon (C) has 6 electrons. It is in the second row, and therefore has two orbitals. Carbon's first two electrons are in the inner orbital and the last four are in the valence electron shell. If you count from left to right in period 2, carbon is the fourth element. This corresponds to the number of valence electrons that carbon has, four. Valence electrons are primarily responsible for an element’s chemical reactivity.

Ionic bonds are extremely strong bonds between atoms with a highly unequal number of valence electrons. Ionic bonds are formed when charged atoms, or ions, are attracted when one gives up one or more of its electrons to the other atom. An ion is an atom or a molecule in which the total number of electrons is not equal to the total number of protons, giving the atom or molecule a net positive or negative electrical charge. Elements that freely give up electrons are known as cations, which have a positive charge due to the loss of an negatively charged electron (-1 x -1 = +1). Cations are elements in Group 1 and Group 2. Elements in these groups have their valence shell nearly complete have a strong negative charge and steal cation’s electrons (+1 x -1 = -1). Such elements are known as anions, and are typically found in Group 16 and Group 17. Ionic bonds occur due to the electrostatic force of attraction between two oppositely charged ions: cations (+) and anions (-).

Table salt is a classic example of an ionic bond. Chemically, salt is known as sodium chloride and has the chemical formula, NaCl. The chemical formula tells us that one atom of sodium (Na) combines with one atom of chlorine (Cl). Let’s see how this bond is formed.

Figure 4. Ionic bond between sodium (Na) and chlorine (Cl) to from sodium chloride, Na+Cl-. The loss of a negatively charged electron (e-) of a sodium atom creates a positively charged sodium cation (Na+), whereas the chlorine becomes a negatively charged anion (Cl-) due to the gain of an electron.

Below is an example of an animated visualization how ionic bonding works in sodium fluoride Na+F-. Like chlorine, fluorine also has one valence electron. So the ionic bond forms similarly to Na+Cl-.

Figure 5. Ionic bonding of NaF. Sodium (Na) gives up one electron to fluorine (F). Since Na lost one negatively charged electron (e-), it has a charge of +1 (-1 x -1 = +1). Fluorine has a charge of -1 since it gained an electron.

Exercsie C: Lewis Dot Diagrams

Chemical reactivity is predominately affected by the number of valence electron atoms have. The electrons of inner orbitals tend to be extremely stable and unreactive. In contrast, the electrons in the outer orbital (the valence orbital), actively react with other atoms. Lewis dot diagrams were developed as a way to simplify the Bohr model of the atom to more efficiently visualize interactions of the valence electrons.

Figure 6. Lewis dot diagram of nitrogen. Nitrogen has five valence electrons. Two valence electrons will pair with each other and be unreactive in a covalent bond. The other three unpaired valence electrons will react with other atoms in a covalent bond.

To construct a Lewis dot diagram, determine the number of valence electron an atom has. For example, nitrogen (N) has five valence electrons. You can easily determine this by locating nitrogen on the periodic table (atomic number 7). In the first three periods of the table, simply count from left to right in the period in which nitrogen is located. Nitrogen is the fifth element in the second period. Therefore, nitrogen has five valence electrons.

For elements in the first three periods (and all elements in groups 1-2 & 13-18), electrons in an orbital tend to pair up when there are more than four valence electrons. In the example of nitrogen, two of the electrons pair up and are unreactive, leaving three electrons to react with neighboring atoms. For the Lewis dot diagram, we diagram this phenomenon by pairing two electrons on one side of the chemical symbol (indicating non-reactivity of those electrons) and placing the other three reactive electrons on the other three sides of the chemical symbol.

Exercsie D: Covalent Bonds

Covalent bonds involves the sharing of electron pairs between atoms, when there is relatively equal attraction due to similar electronegativity between atoms. This relative equality in electromagnetic attraction allows these atoms to share electrons between atoms, bonding them together. Atoms in covalent molecules share enough electrons to complete their valence shell.

Figure 7. Using a Lewis dot diagram to visualize how water (H2O) is formed by covalent bonding. The two molecules of hydrogen have one reactive electron in their valence orbital and each share an electron with a single oxygen atom. Oxygen has six valence electrons, two pairs which are unreactive. The remaining two electrons react with the hydrogen atoms forming single covalent bonds between the oxygen atom and hydrogens.

Figure 8. Skeletal structure of water. Single bonds, in which two electrons are shared between atoms are visualized as a single line between the atoms.

Single bonds

Since hydrogen only has one electron in the first valence shell, it only needs to share one electron from another atom to complete its valence shell. The first shell can only fit two electrons. Elements that covalently bond in periods 2 and 3 (groups 13-16) need to share enough electrons to fill their valence shells, which is eight electrons. The atom of oxygen by itself has 6 valence electrons and needs to share two electrons from neighboring atoms. Each hydrogen atom shares its electron with one of the oxygen’s electrons, pairing up, creating a single covalent bond, typically called a single bond. In a Lewis dot structure, this bond is visualized by the two dots between the H and the O. Oxygen’s four electrons are paired up and are unreactive, which is visualized by the two dots above and below the O. The Lewis dot structure can be simplified into a skeletal structure by highlighting the bonds between the atoms in the molecule, with one line for each two electrons shared of a single bond. The unreactive electrons can be shown in a skeletal structure, but typically they are omitted (yet inferred).

Double bonding

Some molecules can share more than two electrons between atoms. When this happens, double or triple bonds emerge. Oxygen (O2) is an example of a double bond. Each oxygen atom needs to share two of its electrons in order to fill its valence electron. If each atom of oxygen shares its two electrons with another oxygen’s two electrons (four total reactive electrons shared), collectively their valence shells are complete. This is an example of a double bond. The skeletal structure of O2 (Fig. 9) has two lines between the oxygen atoms, representing two pairs of electrons being shared between the two oxygen atoms. A Lewis dot structure of O2 would represent a double bond with four dots between the oxygen atoms, representing the four shared electrons.

Figure 9. Lewis structure and skeletal structure of O2. a) The Lewis structure of O2. An oxygen atom has 6 valence electrons. In O2, two pairs of valence electrons are shared between the two atoms of oxygen, forming a double bond. In a Lewis structure, the double bond is visualized as four dots between the two oxygens. Each oxygen has two pairs of unreactive electrons, shown as dots that are not adjacent between the two atoms. b) The Skeletal structure of O2. O2 is typically visualized as a skeletal structure, with two Os connected by two lines, representing the double bond. Customarily the unreactive electron pairs are not shown.

Exercsie E: Hydrogen bonds

Figure 10. Three dimensional structure of a water molecule. Water is a polar molecule, due to an uneven distribution of electron density. The oxygen in water has a partial negative charge (δ-) due the unequal sharing between the pairs of electrons of oxygen and hydrogen, with oxygen holding the shared electrons more than the hydrogen. In contrast, hydrogen has a partial positive charges (δ+). These partial charges within a water molecule are responsible for hydrogen bonding between water molecules and other molecules, as well.

Water (H2O) is actually a bent molecule. Oxygen has two lone pairs of electrons that do not bond with another atom. Since water is a tetrahedral molecule (think of a styrofoam ball with four toothpicks emerging from it equally distant from each other), the lone electron pairs must be adjacent to each other. When this happens, the lone pairs exert a repulsive force on each other pushing against each other at 121.5˚. Each lone electron pair is also 121.5˚ from its adjacent molecule. This bent nature generates polarity, and bent structure, in the water molecule.

Electronegativity

While a water molecule is formed by covalent bonds, there is unequal sharing of electrons between the oxygen atom and hydrogen atoms. Since oxygen has six valence electrons and hydrogen only has one, oxygen has a stronger pull on the electrons than hydrogen. This pull of an atoms is known as electronegativity.

Oxygen atoms are highly electronegative which attracts electrons to it more strongly than hydrogen, making the region near the oxygen slightly more negative than areas around the two hydrogen atoms. Effectively this makes oxygen partially negative (δ-). In other words, the oxygen atom holds electrons more often than hydrogen. Hydrogen’s low electronegativity is due to its relative inability to hold onto its electrons in the presence of oxygen, gives it a partially positive (δ+) charge. This difference in polarity of water gives its many unique properties. Hydrogen bonding happens between hydrogen and the following highly electronegative elements: oxygen (O), nitrogen (N), and fluorine (F).

Hydrogen bonding in water

In water, hydrogen bonds form between the partially negative (δ-) oxygen atoms of one water molecule with the partially positive (δ+) hydrogen atoms of another water molecule. In liquid water, hydrogen bonding occurs between molecules, but hydrogen bonds are easily broken because they are weak. Therefore, water molecules are continually bonding and detaching. In ice, water molecules bond with neighboring molecules, but they do not detach. Due to the bent structure of the water molecule, the molecular structure of ice forms a lattice network that is more spread out than liquid water.

Figure 11. Hydrogen bonding in liquid water. The electronegativity of a water molecule is responsible for the liquid nature of water. Hydrogen bonds form between adjacent O atoms and H atoms of different H20 molecules. These molecules readily form and detach. Liquid water molecules are randomly assorted and continually attaching and detaching, giving the property of liquid water a fluid state. Liquid water is much denser than ice, because its molecules are more compact (dense) than the crystalline lattice of ice.

Figure 12. Hydrogen bonding in solid water, or ice. Ice is a solid at a molecular level because hydrogen bonds form among all the water molecules forming a lattice structure. This lattice structure (or crystalline structure) is more spread out than liquid water, due to the polarity (and bent shape) of the water molecule. Molecules that are more spread out generate substances that have lower density. This is why ice floats on water.


Teacher Notes: Chemical Bonds and Forces

Intramolecular bonds are the bonds that hold atoms to atoms and make compounds. There are 3 types of intramolecular bonds: covalent, ionic, and metallic.

Covalent Bond: a bond in which a pair or pairs of electrons is shared by two atoms.

  • Molecular compounds refer to covalently-bonded species, generally of low molecular mass.
  • Macromolecular compounds are high molecular mass compounds that are covalently-bonded and linear, branched, or cross linked.
  • Network: compounds in which each atom is covalently-bonded to all its nearest neighbors so that the entire crystal is one molecule.

Ionic Bond: a bond that holds atoms together in a compound the electrostatic attraction between charged ions. Ionic compounds are formed between atoms that differ significantly in electronegativity. The electron(s) involved in bonding is (are) transferred from the less electronegative to the more electronegative atom(s) forming ions.

Metallic Bond: a bond resulting from the attraction between positive ions and surrounding mobile electrons.

Intermolecular forces

Intermolecular forces are the forces that attract molecules or particles to like or unlike molecules or particles. Typically, these forces between molecules form much weaker bonds than those bonds that form compounds. Intermolecular forces are described below. They are grouped into 3 subcategories based on the type of intramolecular bonds that form a compound:

  • Ionic compounds exhibit electrostatic intermolecular forces that form strong bonds with other ionic species.
  • Covalent compounds exhibit van der Waals intermolecular forces that form bonds of various strengths with other covalent compounds. The three types of van der Waals forces include: 1) dispersion (weak), 2) dipole-dipole (medium), and 3) hydrogen (strong).
  • Ion-dipole bonds (ionic species to covalent molecules) are formed between ions and polar molecules. These compounds typically form medium to strong bonds.

There are five kinds of intermolecular forces described below the bond strengths described range from strongest to weakest (the latter 3 are examples of van der Waals forces). Please remember that this comparison is relative to other intermolecular attractions and not to covalent or ionic bond strength there are numerous exceptions that are not provided here.


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Some Books about Biology and Evolution

Futuyma, Douglas J. (1997). Evolutionary Biology. Sunderland, Mass.: Sinauer Associates.

Ridley, Mark. (2003). Evolution. Boston: Blackwell Scientific.

Hartl, Daniel L. & Andrew G. Clark. (1997). Principles of Population Genetics. Sunderland, Mass.: Sinauer Associates.

Crow, James F. & Motoo Kimura. (1970). Introduction to Population Genetics Theory. Edina, Minn.: Burgess Publishing Company.

Graur, Dan & Wen-Hsiung Li. (2000). Fundamentals of Molecular Evolution. Sunderland, Mass.: Sinauer Associates.

Lewontin, Richard C. (1974). The Genetic Basis of Evolutionary Change. New York: Columbia Univ. Press.

Gillespie, John H. (1997). The Causes of Molecular Evolution. New York: Oxford Univ. Press.

Golding, Brian, ed. (1994). Non-Neutral Evolution. Boston: Chapman and Hall.

Endler, John A. (1986). Natural Selection in the Wild. Princeton, N.J.: Princeton Univ. Press.

Cowen, Richard. (2004). History of Life. Boston: Blackwell Scientific.

Dawkins, Richard. (1987). The Blind Watchmaker. New York: W.W. Norton.

Kitcher, Philip. (1982). Abusing Science. Cambridge, Mass.: MIT Press.

Wilson, Edward O. (1992). The Diversity of Life. Cambridge, Mass.: Harvard Belknap.

Haldane, J.B.S. (1932). The Causes of Evolution. Princeton, N.J.: Princeton Univ. Press (reprinted 1990).

Simpson, George G. (1944). Tempo and Mode in Evolution. New York: Columbia Univ. Press.

Mayr, Ernst E. (1982). The Growth of Biological Thought. Cambridge, Mass: Harvard Belknap.


Acknowledgements

Tet TKO mESC lines were kindly provided by G.-L. Xu (Shanghai Institutes for Biological Sciences) and R. Jaenisch (Whitehead Institute, MIT, Cambridge). We are grateful to M. Okano and H. Niwa (both at Kumamoto University, Japan) for providing the Dnmt TKO mESC line and the Oct4-YFP reporter cell line, respectively. A.S.S. is supported by a fellowship from the Fonds der Chemischen Industrie. We thank the Deutsche Forschungsgemeinschaft for financial support through the programs: SFB749 (TP A4), SFB1032 (TP A5), SPP1784 and CA275-11/1. We thank the European Union Horizon 2020 program for funding the ERC Advanced project EPiR (741912). Further support is acknowledged from the Excellence Cluster CiPSM (Center for Integrated Protein Science).


Biology Notes on Enzymes

Here is a compilation of notes on enzymes. After reading these notes you will learn about: 1. Introduction to Enzymes 2. Origin of Enzymes 3. Historical Landmarks 4. Meaning 5. Importance 6. Unit 7. Chemical Nature 8. Properties 9. Characteristics 10. Nomenclature 11. Classification 12. Enzymes Vs. Non-Biological Catalysts 13. Catalysts and Enzymes 14. Types 15. Modes of Enzyme Action and Others.

  1. Notes on Introduction to Enzymes
  2. Notes on the Origin of Enzymes
  3. Notes on the Historical Landmarks of Enzymes
  4. Notes on the Meaning of Enzyme
  5. Notes on the Importance of Enzyme
  6. Notes on the Unit of Enzyme
  7. Notes on the Chemical Nature of Enzymes
  8. Notes on the Properties of Enzyme
  9. Notes on the Characteristics of Enzymes
  10. Notes on the Nomenclature of Enzymes
  11. Notes on the Classification of Enzymes
  12. Notes on Enzymes Vs. Non-Biological Catalysts
  13. Notes on Catalysts and Enzymes
  14. Notes on the Types of Enzymes
  15. Notes on the Modes of Enzyme Action
  16. Notes on the Inhibition of Enzyme Action
  17. Notes on the Feedback Inhibition of Enzymes
  18. Notes on the Specificity of Enzyme
  19. Notes on Factors Influencing Enzyme Activities
  20. Notes on the Biochemical Pathways of Enzymes
  21. Notes on the Regulation of Enzyme

Note # 1. Introduction to Enzymes:

Thousands of chemical reactions proceed very rapidly at any given instant within all living cells of an organism. Virtually all of these reactions are mediated by remarkable molecular devices called enzymes. That is, the enzymes are central to every biochemical reaction and are called the catalysts of biological systems (biocatalysts).

They in organized sequences and catalyse the hundreds of stepwise reactions by which nutrient molecules are degraded, chemical energy is conserved and transformed, and biological macromolecules are made from simple precursors. Through the action of regulatory enzymes, metabolic pathways are highly coordinated to yield a harmonious interplay among the many different activities necessary to sustain life.

Enzymes catalyse an enormous diversity of biochemical reactions due to their capacity to specifically bind a very wide range of molecules. By utilizing the full repertoire of intermolecular forces, enzymes bring substrates together in an optimal orientation, the prelude to making and breaking chemical bonds.

They catalyse reactions by stabilizing transition states, the highest energy-species in reaction pathways. By selectively stabilizing a transition state, an enzyme determines which one of several potential biochemical reactions actually takes place.

Until 1980s, all enzymes were believed to be proteins. Then, Tom Cech and Sidney Altman independently discovered that certain RNA molecules may function as enzymes may be effective biocatalysts. These RNA biocatalysts have come to be known as ribozymes.

Note # 2. Origin of Enzymes:

Enzymes are commonly proteinaceous substances which are capable of catalysing chemi­cal reactions of biological origin without themselves undergoing any change. Therefore, they are called biocatalysts. Enzymes are synthesised by living cells.

The term ‘enzyme’ was coined by Kuhne (1878) for catalytically active substances previously called ferments. Enzymes were actually found out by Buchner (1897) with the accidental discovery that fermentation of sugar is not only caused by living yeast cells but also yeast extract.

The extract obviously possessed biocatalysts required for the process. Buchner (1903) also isolated the first enzyme. He was awarded Nobel Prize in the same year, 1903. There are numerous enzymes as every biochemical reaction is catalysed by a separate enzyme. It is estimated that a cell contains over 5000 chemicals. The number of chemical reactions is many times more.

Therefore, the number of enzymes is several thousands. A cell with an average diameter of 20 pm has about 1000 chemical reactions going on at any time. All of them require specific enzymes. All the enzymes are not present at all the times in the cell but they are formed as and when required from the blue print present in DNA.

Enzymes are mainly functional inside the living cells. As found out by Buchner they can be extracted from the cells and made to catalyse reactions outside the living cells. In nature some enzymes are secreted by living cells and made to perform extracellular catalysis.

Digestive enzymes belong to this category. Several enzymes of medical and chemical impor­tance are now available in the market, e.g., rennet tablets (from rennin of calf stomach) for coagulating milk protein casein during preparation of cheese and other milk products.

En­zymes functional outside the living cells are called exo-enzymes, e.g., enzymes present in digestive juices, lysozyme of tears. Enzymes functional inside living cells are known as endoenzymes, e.g., enzymes of Krebs cycle (inside mitochondria), enzymes of glycolysis (inside cytoplasm).

The biochemical which is acted upon by an enzyme is known as substrate. In case two bio-chemicals are involved in a reaction, the same are called reactants. The chemicals formed after the completion of a reaction are termed as products. The final products are also called end products. Part of enzyme that takes part in catalysing biochemical reaction is called active site.

Note # 3. Historical Landmarks of Enzymes:

The very existence of biological catalysis was first recognized and described during the late 18th century while studying the digestion of meat by secretion of the stomach. Subsequently, Louis Pasteur concluded in 1850s that fermentation of sugar into alcohol by yeast is catalyzed by “ferments”.

He postulated that these ferments were inseparable from the structure of living yeast cells, a view called ‘vitalism’ that prevailed for many years. F.W. Kuhne coined the term enzyme in 1878 to represent the “ferments”. The first enzyme was isolated by E. Buchner in 1903 for which he was awarded Nobel Prize in the same year.

The protein nature of enzyme was first discovered by James Sumner in 1926 when he purified the enzyme urease and obtained it in crystalline form. For this, Sumner was awarded Nobel Prize in 1946.

Note # 4. Meaning of Enzyme:

An enzyme is a protein that is synthesised in a living cell and catalyses or speeds up a thermodynamically possible reaction so that the rate of the reaction is compatible with the biochemical process essential for the maintenance of the cell. It is sometimes called as organic catalyst or biocatalyst.

Over 90% of enzymes are simple globu­lar proteins (Fig. 8.14). The remainder is conjugated proteins, which have a non­-protein fraction called the prosthetic group. Many enzymes have relative molecular mass of between 10,000 and 50,000da.

The first enzyme discovered was amy­lase, which catalyses the conversion of starch to maltose, in 1833 by two French chemists Payen and Persoz. However, it was not well-known until 1876 when Wilhelm Kuhne, the distinguished German biochemist, proposed the term enzyme.

Note # 5. Importance of Enzyme:

Biological Importance of Enzymes:

(i) Thousands of chemical reactions are taking place in the body of a living organism. All of them are mediated by enzymes,

(ii) Enzymes are specialised catalysts that operate at biological temperatures,

(iii) Enzyme mediated reactions do not require harsh treatment,

(iv) They are pH specific so that reactions requiring different pH operate in different parts of the body,

(v) As they operate under favourable conditions, enzymes force the organisms to live under favourable environment,

(vi) Enzymes are highly regulated. Their formation is controlled by separate genes. Activation and repression of genes allow certain enzymes to be functional or non-functional in cells.

Economic Importance of Enzymes:

It is enzyme based detection of diseases like AIDS.

They are enzymes used in breaking DNA at specific sites. DNA frag­ments are employed in genetic engineering.

iii. Alcoholic Beverages:

Enzyme complex zymase obtained from yeast is used in brewing or fermentation of alcoholic drinks.

They contain protease for brighter washing of clothes and amylase for dish washing.

Trypsin is added to partially pre-digest baby foods.

The enzyme is used in clearing blood clots inside blood vessels.

Diastase and other enzymes are used regularly by patients with deficient digestive juices.

Rennet or rennin tablets are used for preparation of cheese. Lactase and lipase are employed to provide proper consistency and flavour to cheese.

It is used for clearing fruit juices, retting of fibres and preparation of green coffee.

The enzyme is employed for chill-proofing of beverages, degumming of silk, cleaning of hides, softening of bread and meat.

Note # 6. Unit of Enzyme:

The actual molar amount of the enzyme in an enzyme-catalyzed reaction is not known in many situations. In such cases the amount of enzyme can be expressed in terms of the enzyme activity observed.

The International Commission on Enzymes established by International Union of Biochemistry defines One International Unit of enzyme as the amount of enzyme that catalyzes the formation of one micromole of product in one minute.

In determining the One International Unit the conditions of assay must be specified because enzymes are very sensitive to factors such as pH, temperature, and ionic strength. Another definition for units of enzyme is the ‘katal’. One katal is defined as the amount of enzyme that catalyses the conversion of one mole of substrate to product in one second. Thus, one katal equals 6 x 107 international units.

Note # 7. Chemical Nature of Enzymes:

All enzymes are globular proteins with the exception of recently discov­ered RNA enzymes. Some enzymes may additionally contain a non-protein group. Accordingly there are two types of enzymes, simple and conjugate.

It is an enzyme which is wholly made up of protein. Active site is formed by specific grouping of its own amino acids. Additional substance or group is absent, e.g., pepsin, trypsin, urease.

It is an enzyme which is formed of two parts— a protein part called apoenzyme (e.g., flavoprotein) and a nonprotein part named cofactor. The complete con­jugate enzyme, consisting of an apoenzyme and a cofactor, is called holoenzyme. Active site is formed jointly by apoenzyme and cofactor.

Cofactor is small, heat stable and dialysable part of conjugate enzyme. It may be inor­ganic or organic in nature. Organic cofactors are of two types, coenzymes and prosthetic groups.

Coenzymes are easily separable non-protein organic cofactors. Prosthetic groups are non-protein organic cofactors firmly attached to apoenzymes, e.g., heme (=haem), biotin, pyridoxal phosphate. Heme (= haem) is iron containing prosthetic group in cytochromes, haemoglobin, myoglobin, catalase and peroxidase.

The last two cause breakdown of hydro­gen peroxide to water and oxygen. FMN and FAD are considered prosthetic groups by some workers while others consider them to be coenzymes.

Both coenzyme and prosthetic group take part in group transfer reactions. Prosthetic group requires a single apoenzyme for picking up the group and transferring the same. Coenzyme requires two Apo enzymes, one for picking up the group and the second for transferring the group, e.g., NAD + , NADP + , CoA.

Coenzyme has three important functions:

(a) Coenzyme is essential for bringing the substrate in contact with the enzyme,

(b) It picks up a product of the reaction, e.g., hydrogen in case of NAD + (nicotinamide adenine dinucleotide) or NADP + .

(c) The product picked up by a coenzyme is transferred to another reactant.

Certain workers use the term cofactor for any loosely bound non-protein group. The organic cofactor is called coenzyme. They use the term prosthetic group similarly for both inorganic and organic group attached firmly to apoenzyme.

Most of the coenzymes are made of water soluble vitamins, В and C, e.g., thiamine, riboflavin, nicotinamide, pyridoxine. Inorganic cofactors include ions of a variety of min­erals e.g., calcium, iron, copper, zinc, magnesium, manganese, potassium, nickel, molybde­num, selenium, cobalt.

They usually function as activators by forming one or more coor­dination bonds with both the substrate and active site of enzyme. Fe 2+ is cofactor for catalase. Chloride ion stimulates activity of salivary amylase. Zinc is required for carboxypeptidase NAD + and NADP + activity.

Active Site or Active Spot:

The whole of enzyme molecule is not active in catalysing a chemical reaction. Only a small portion of it is active. It is called active site or active spot. An enzyme may have one to several active sites. An active site or spot is an area of the enzyme which is capable of attracting and holding particular substrate molecules by its specific charge, size and shape so as to allow the chemical change.

It fails to recognise other molecules. Active site consists of a few amino acids and their side groups which are brought together in a particular fashion due to secondary and tertiary folding of a protein molecule (Fig. 9.27) and its association with the cofactor, if any.

For example, the active site for aldolase is glycine-histidine-alanine while that of pyruvic oxidase is aspartic acid-cysteine-alanine. The remaining amino acids help maintain the shape of the enzyme molecule.

Note # 8. Properties of Enzyme:

Some of the properties of enzyme are as follows:

Enzymes are generally globular proteins. They may have additional inorganic or organic substances for their activity. However, two types of RNA enzymes are known, ribozyme and ribonuclease-P. Peptidyl transferase has also been found to be part of rRNA by Noller (1992).

ii. Molecular Weight:

Being proteinaceous, the enzymes are giant molecules with a molecular weight of 6000 (bacterial ferredoxin) to 4,600,000 (pyruvate dehydrogenase complex).

iii. Colloidal Nature:

They are hydrophilic and form hydrosol in the Free State.

iv. Chemical Reaction:

Enzymes do not start a chemical reaction but increase the rate of chemical reaction. They do not change the equilibrium but bring about equilibrium very soon.

v. Efficiency:

The number of substrate molecules changed per minute by a molecule or enzyme is called turn over number (kcat). The higher the turn-over number, the more efficient an enzyme is. It depends upon the number of active points present over an enzyme, precise collisions between reactants and the rate of removal of end products.

The optimum turn-over number for enzyme carbonic anhydrase (enzyme present in RBCs) is 36 million, catalase 5 million, enzyme sucrase or invertase 10,000 and flavoprotein 50. Enzyme effi­ciency is usually much more than that of inorganic catalysts.

For example:

The rate of enzyme mediated hydrolysis of urea is 1014 times higher than the rate of its acid hydrolysis carried out at 40°C higher temperature. Similarly, the rate of CO2 hydration is 10 million times faster in the presence of enzyme carbonic anhydrase than in its absence.

vi. Unchanged Form:

Enzymes are in no way transformed or used up in the chemical reaction but come out unchanged at the end of reaction.

vii. Reversibility:

Theoretically, all enzyme controlled reactions are reversible. Reversibility is, however, dependent upon energy requirements, availability of reactants, concentration of end products and pH.

viii. Enzyme Specificity:

Enzymes are highly specific in their action. For example, en­zyme maltase acts on sugar maltose but not on lactose or sucrose. Different enzymes may act on the same substrate but give rise to different products.

For example, raffinose gives rise to rnelibiose and fructose in the presence of enzyme sucrase while in the presence of enzyme melibiase it produces lactose and sucrose. Similarly an enzyme may act on different substrates, e.g., sucrase can act on both sucrose and raffinose producing different end products.

ix. Heat Sensitivity:

All enzymes are heat sensitive or thermolabile. Most enzymes operate optimally between 25°-35°C. They become inactive at freezing temperatures and denatured at 50°-55°C. However, thermal algae and bacteria are an exception. Their enzymes remain functional even at 80°C. Enzymes of seeds and spores are also not denatured at 60°— 70°C.

x. Protein Poisons:

Being made of proteins, enzymes are inactivated or de­natured by all those substances and forces which destroy protein structure, e.g., heavy metals, high energy radia­tions.

Each enzyme functions at a particular pH (Fig. 9.28.), e.g., pepsin (2 pH), sucrase (4.5 pH), salivary amy­lase (6.8 pH), trypsin (8.5 pH). A change in pH makes the enzymes ineffective.

Specificity of pH for enzyme ac­tivity is useful in regulating enzymes, e.g., salivary amylase stops its activity in stomach where hydrochloric acid is secreted. The same acid activates another enzyme pepsin from its pre­cursor called pepsinogen. Pepsinogen can also be changed into pepsin by catalytic activity of the latter.

xii. Enzyme-Substrate Complex:

The active sites of enzymes have a specific conformation for attracting and holding substrate. It usually possesses a crevice or pocket where the substrate fits in a complementary fashion. The two join to form a complex known as enzyme—substrate complex (ES).

The complexed state is short lived. The substrate is changed into products. The products remain complexed with the active site of the enzyme for a brief period. They soon separate and the active site is freed to perform another catalytic act.

The greater the affinity of the enzyme for a substrate, the higher is the catalytic activity.

xiii. Chain Reactions:

Biochemical reactions are not isolated. A number of them occur in quick succession. A team of enzymes work one after the other to accomplish such multistep reactions, e.g., five enzymes for conversion of threonine to isoleucine.

Note # 9. Characteristics of Enzymes:

All enzymes are proteins, but a func­tional enzyme has different components and these components are named differently, viz.,

A conjugated protein and functional enzyme.

Polypeptide segment of the enzyme, which is catalytically inactive.

The non-protein organic moiety, which can frequently be separated from the apoenzyme.

If a substance is firmly (covalently) attached to the protein part of the enzyme, it is referred to as a prosthetic group. It is the non-protein portion of any conjugated protein. So coenzyme is a specific example of prosthetic group.

There are many metalloprotein enzymes in which the metal ion (e.g. Mg ++ , Mn ++ , and Zn ++ ) is bonded either to the apoen­zyme or to the coenzyme. The metal is usu­ally designated as activator. They form a co-ordination complex between the enzyme and the substrate, and activate the substrate by prompting electronic shifts.

Pro-enzyme or Zymogens:

They are simple protein enzymes, which are secreted, in an inactive form.

It is the process in which an inactive protein (pro-enzyme or zymogens) is transformed into an active enzyme.

Note # 10. Nomenclature of Enzymes:

All enzyme names should end in suffixase. Exceptions are some old names, e.g., ptyalin, pepsin, trypsin. Some old names indicate the source but not the action, e.g., papain from Papaya, bromelain from Pineapple of family Bromeliaceous.

In modem system enzyme names are given after:

(i) Substrate acted upon, e.g., sucrase (after sucrose), lipase, proteinase, nuclease, peptidases, maltase

(ii) Chemical reaction, e.g., dehydrogenase, oxi­dase, carboxylase, decarboxylase, etc.

The second category of names are group names. They are often qualified by the addition of the name of substrate, e.g., succinic dehydroge­nase, isocitric dehydrogenase, glutamate-pyruvate transaminase, DNA polymerase.

Thus DNA polymerase catalyses synthesis of DNA segments through polymerisation of deoxyribonucleotides. Similarly glutamate-pyruvate transaminase transfers amino group (—NH2) from glutamate to pyruvate.

Note # 11. Classification of Enzymes:

In older times enzymes were classified into two broad categories:

(i) Hydrolysing:

Catalysing hydrolysis of larger molecules into smaller ones, e.g., carbohydrates or amylases, proteases, lipases, esterases, phosphorylases, amidases. Digestive enzymes are hydrolysing in nature. They are often grouped into three types— proteolytic, amylolytic and lipolytic,

(ii) Desmolysing:

Catalysing reactions other than hyrolysis, e.g., aldolases, dehydrogenases, oxidases, peroxidases, catalases, carboxylases, etc. The modem system of enzyme classification was introduced by International Union of Biochemistry (IUB) in 1961. It groups enzymes into the following six categories.

They take part in oxidation and reduction reactions or transfer of electrons.

Oxidoreductases are of three types— oxidases, dehydrogenases and reductases, e.g., cytochrome oxidase (oxidises cytochrome), succinate dehydrogenase, nitrate reductase.

They transfer a group from one molecule to another e.g., glutamate- pyruvate transaminase (transfers amino group from glutamate to pyruvate during synthesis of alanine). The chemical group transfer does not occur in the Free State.

They catalyse hydrolysis of bonds like ester, ether, peptide, glycosidic, С-С, С halide, P—N, etc. which are formed by dehydration condensation. Hydrolases break up large molecules into smaller ones with the help of hydrogen and hydroxyl groups of water molecules. The phenomenon is called hydrolysis. Digestive enzymes belong to this group, e.g., amylase (hydrolysis of starch), sucrase, lactase.

The enzymes cause cleavage, removal of groups without hydrolysis, addi­tion of groups to double bonds or removal of a group producing double bond, e.g., histidine decarboxylase (breaks histidine to histamine and CO2), aldolase (fructose-1, 6-diphosphate to dihydroxy acetone phosphate and glyceraldehyde phosphate).

Fructose 1, 6-diphosphate – aldolase → Dihydroxy acetone phosphate + Glyceraldehyde phosphate.

The enzymes cause rearrangement of molecular structure to effect isomeric changes. They are of three types, isomerases (aldose to ketose group or vice-versa like glucose 6-phosphate to fructose 6-phosphate), epimerases (change in position of one constituent or carbon group like xylulose phosphate to ribulose phosphate) and mutases (shifting the position of side group like glucose-6-phosphate to glucose-1- phosphate).

f. Ligases (Synthetizes):

The enzymes catalyse bonding of two chemicals with the help of energy obtained from ATP resulting in formation of such bonds as С-О, С-S, С-N and P-O, e.g., pyruvate carboxylase. It combines pyruvic acid with CO2 to produce oxaloacetic acid.

Most of the chemical reactions do not start automatically because the reactant molecules have an energy barrier to become reactive.

The energy barrier may be on account of:

(i) Mutual repulsion due to presence of electrons over their surfaces,

(ii) Solvation or holding of reactants in solution form by hydrogen bonds,

(iii) Reaction sites of the reactive molecules being small, precise collisions do not occur.

Therefore, an external supply of energy is needed for the start of the chemical reaction. It is called activation energy. Activation energy increases the kinetic energy of the system and brings about forceful collisions between the reactants. The requirements of activation energy is quite high. For example, acidic hydrolysis of sucrose requires 32000 cal/ mole of energy.

As already noted about 1000 chemical reactions are taking place in a cell at any time. Activation energy required for such a large number of reactions cannot be provided by living systems. Enzymes lower the activation energy required for a reaction (Fig. 9.32).

For example, in the presence of enzyme sucrase or invertase, hydrolysis of sucrose requires an activation energy of 9000 cal/mole (instead of 32,000 cal/ mole).

This is achieved by four ways:

(i) De-solvation or taking the reactants out of solution state,

(ii) Establishing weak bonds between reactants and enzyme. It releases energy called bond energy,

(iii) Bringing the reactant molecules closer in the region of active sites of enzymes,

(iv) Development of strain in the bonds of the reactants by electrophilic and nucleophilic attack,

(v) Formation of unstable intermediate structural states collectively called transition state. During the transition state the substrate bonds are broken and new bonds are established that transform the substrate molecules into products,

(vi) In exother­mic reactions, the energy content of the products is lower than that of substrate (Fig. 9.32).

It is higher in case of endothermic reactions. However, whether the reaction is endothermic or exothermic, energy is required for pushing the substrate molecules into transition state. The difference in the energy level of substrate (S) and transition state is the activation energy required to start the reaction.

Note # 12. Enzymes Vs. Non-Biological Catalysts:

A catalyst is a molecule that accelerates a particular chemical reaction without itself being chemically altered in the process, and may be biological or non-biological in origin. Biological catalysts are the enzymes.

Enzymes are similar to non-biological catalysts in the following respects:

(i) They lower the activation energy of reaction,

(ii) They do not participate in the reaction, and return in their original form at the end of reaction, and

(iii) They only increase the reaction rate.

But enzymes increase the rate of reaction at a phenomenal scale, and are highly specific. These features are unimaginable for non-biological catalysts (Table 27.1). Many enzymes may be less specific in binding to the substrate, but they are always extremely specific in the reaction they catalyze. For example, mammalian cytochrome P450 binds to a variety of substrates, but it always adds a -OH group to the substrate.

But many enzymes are highly specific in binding as well, e.g., glucose oxidase binds only to D- glucose. Enzymes can distinguish between similar parts of the substrate molecule (this property is called regiospecificity) and between optical isomers of the substrate (this ability is called stereo-specificity).

In addition, enzymes are subject to a variety of regulations, and their reaction rates show substrate saturation, which is not the case with catalysis (non-biological).

Enzymes are attractive because they operate under mild conditions of temperature, pressure and pH, which saves energy, and undesirable by-products are not produced by enzymes this simplifies product recovery. Finally, certain stereo-chemical reactions are impractical with chemical methods.

The disadvantages associated with enzymes are as follows:

(i) High costs of enzymes, and

(ii) General instability of purified enzymes so much so that some enzymes cannot be used due to instability.

Note # 13. Catalysts and Enzymes:

Catalysts are inorganic substances which increase the rate of chemical reactions without themselves undergoing any change and without modifying the equilibrium of the reactions. Enzymes are similar chemicals which are biological in origin and operate in the biochemical world.

Both catalysts and enzymes remain unchanged chemically and quanti­tatively at the end of the reaction, so that they can be used over and over again.

They are required in minute quantity as compared to their substrate.

Reactions controlled by both catalysts and enzymes are theoretically reversible though reversibility is dependent upon different kinetics.

They do not change the equilibrium of the reaction.

Both catalysts and enzymes increase the rate of chemical reac­tion. They do not initiate the reaction.

They lower the activation energy required for starting the chemi­cal reaction.

They form short lived complexes with the substrate molecules.

End products of a reaction are not changed by catalysts and enzymes.

Note # 14. Types of Enzymes:

Enzymes are of three types:

i. Pro-Enzyme or Zymogen:

Pro-enzyme is the inactive precursor of an enzyme. The term zymogen is often used for inactive precursor of proteolysis enzyme, e.g., pepsinogen for enzyme pepsin. Many en­zymes are initially produced in the pro-enzyme or zymogen state.

They become reactive or active enzymes only at a particular pH, in the presence of substrate or some special treat­ment. For example, pepsinogen is changed to active enzyme pepsin in the presence of hydrochloric acid of gastric juice. Thereafter, pepsin has autocatalytic effect on further conversion of pepsinogen.

ii. Allosteric Enzymes:

They are enzymes which have separate areas for different types of modulators that alter the conformation of the active site so as to make it effective or ineffective (Fig. 9.36). The areas are called allosteric sites. The substances which cause change in allosteric sites are known as modulators, allosteric substances or effectors.

The latter are of two types— activators and inhibitors. Allosteric activator binds with an allosteric site in such a way as to make active site operational. Allosteric inhibitor, on the other hand, brings about such a change in the active site that it becomes unable to combine with substrate molecules. For example, the enzyme phosphofructokinase is activated by ADP and inhibited by ATP.

iii. Isoenzymes (Isozymes):

At one time it was believed that an organism has only a single enzyme for a given step of a metabolic reaction. It was later discovered that a substrate may be acted upon by a number of variants of an enzyme producing the same product.

The multiple molecular forms of an enzyme occurring in the same organism and having a similar substrate activity are called isoenzymes or isozymes. Over 100 enzymes are known to have isoenzymes.

Thus a- amylase of wheat endosperm has 16 isozymes, lactic dehydrogenase has 5 isoenzymes in man, while alcohol dehydrogenase has 4 isozymes in maize. Isoenzymes differ in activity optima and inhibition. They are thus useful to organism in adapting to varied environmental conditions.

Note # 15. Modes of Enzyme Action:

There are two view points by which enzymes are supposed to bring about chemical reaction.

i. Lock and Key Hypothesis:

It was put forward by Emil Fischer in 1894. According to this hypothesis, both enzyme and substrate molecules have specific geometrical shapes. ‘In the region of active sites the surface configuration of the enzyme is such as to allow the particular substrate molecules to be held over it. The active sites also contain special groups having —NH2, —COOH, —SH for establishing contact with the substrate molecules.

The contact is such that the substrate molecules or reactants come together causing the chemical change. It is similar to the system or lock and key. Just as a lock can be opened by its specific key, a substrate molecule can be acted upon by a particular enzyme. This also explains the specificity of enzyme action.

After coming in contact with the active site of the enzyme, the substrate molecules or reactants form a complex called enzyme-substrate complex. In the complexed state the molecules of the substrate undergo chemical change.

The products remain attached to the enzyme for some time so that an enzyme-product complex is also formed. However, the products are soon released (Fig. 9.34) and the freed enzyme is able to bind more substrate molecules.

Enzyme + Substrate ⇋ Enzyme – Substrate Complex

Enzyme – Substrate Complex ⇋ Enzyme – Products Complex

Enzyme – Products Complex ⇋ Enzyme + Products

Thus we see that the chemical reactants do not cause any alteration in the composition or physiology of the enzyme. The same enzyme molecule can be used again and again (Fig. 9.35). Hence, enzymes are required in very small concentrations.

1. Blow and Steitz (1970) have found the formation of complex between the enzyme chymotrypsin and its substrate.

2. Keilen and Maun have observed that the absorption spectra of the same enzyme are different in the free state and in the pres­ence of the substrate.

3. The theory explains how a small con­centration of enzyme can act upon a large amount of the substrate.

4. Lock and key theory explains how the enzyme remains unaffected at the end of chemical reaction.

5. It is able to predict the increase in the rate of chemical reaction on the addition of more enzyme or substrate.

6. The theory explains how a substance having a structure similar to the substrate can work as competitive inhibitor.

ii. Induced-Fit Theory (Fig. 9.35):

It is modification of lock and key hypothesis which was proposed by Koshland in 1959. Accord­ing to this theory the active site of the enzyme contains two groups, buttressing and catalytic. The buttressing group is meant for support­ing the substrate. The catalytic group is able to weaken the bonds of reactants by electrophilic and nucleophilic forces.

The two groups are normally at a distance. As soon as the substrate comes in contact with the buttressing group, the active site of the enzyme undergoes conformational changes so as to bring the catalytic group opposite the substrate bonds to be broken.

Catalytic group helps in bringing about chemical reaction. The substrate is converted into product. The product is unable to hold on the buttressing site due to change in its structure and bonds. Buttressing group reverts to its original position. The product is released.

Note # 16. Inhibition of Enzyme Action:

Reduction or stoppage of enzyme activity due to presence of adverse conditions or chemicals is called enzyme inhibition. It is of several types. Inhibition can be classified into two (a) Reversible and irreversible (b) Competitive and non-competitive.

Reversible inhibition is that inhibition which can be overcome by withdrawal of the inhibitor because the effect of the latter is of temporary nature due to blocking of active site or binding to linkages required for maintenance of active site. Dilution and dialysis reduces or eliminates the effect of reversible inhibition. Irreversible inhibition is of permanent nature as the enzyme conformation is harmed.

Denaturation of enzyme is an example of irreversible inhibition. Heavy metals (e.g., Ag + , Hg 2+ , As + ) and iodoacetic acid cause irrevers­ible inhibition by combining with —SH groups and destroying protein structure. Dilution and dialysis have little effect once irreversible inhibition has set in.

Competitive inhibition is caused by swamping of the active sites by a chemical which is similar in structure to the substrate but does not undergo chemical change. Competitive inhibition is usually reversible. Non-competitive inhibition is caused by alteration of con­formation of the enzyme by a chemical that binds to a site other than the active site. It may be reversible or irreversible.

Four common types of enzymes inhibition are as follows:

i. Protein Denaturation:

Enzyme activity is dependent upon the maintenance of ter­tiary structure of the protein moiety. The latter is destroyed by several factors like heat, high energy radiations and salts of heavy metals.

ii. Competitive inhibition:

It is the inhibition of enzyme activity by the presence of a chemical that competes with the substrate for binding to the active site of the enzyme. The inhibitor chemical is also called substrate analogue or competitive inhibitor.

It resembles the substrate in structure and gets bound up to the active site of the enzyme without getting transformed by the latter (Fig. 9.37). As a result, the enzyme cannot participate in catalytic change of the substrate. This is similar to the jamming of a lock by a key similar to original one.

Equilibrium constant for inhibitor binding is called Ki. A high Ki reduces enzyme activity while a low Ki allows enzyme activity to continue though at a reduced rate. Classical example of competitive inhibition is reduction of activity of succinate dehydro­genase by malonate, oxaloacetate and other anions which resemble succinate in their struc­ture.

Competitive inhibition is usually reversible since the addition of more substrate tends to reduce the effect of the inhibitor.

The in­hibition is important in that:

(i) It gives evidence for lock and key hypothesis of enzyme action,

(ii) Substrate analogues are not metabolized by enzymes,

(iii) Control of bacterial pathogens has been effected through competitive inhibition.

Sulpha drugs (e.g., sulphanilamide) inhibit the synthesis of folic acid in bacteria by competing with p-amino benzoic acid (PABA) for the active site of enzyme. Preformed folic acid is obtained by animal cells. Therefore, sulpha drugs do not harm them.

iii. Non-competitive Inhibition:

It is an irreversible inhibition of enzyme activity by the presence of a substance that has no structural similarity with the substrate. It is of two types, reversible and irreversible.

The irreversible non-competitive inhibitor destroys or com­bines irreversibility with a functional group of enzyme that is essential for its catalytic function. Cyanide inhibits the activity of cytochrome oxidase by combining with its metallic ions.

It has no structural similarity with the substrate of the enzyme, namely cytochrome c. Cytochrome oxidase is a respiratory enzyme. In its inhibition, the animal is unable to perform the respiration properly and gets killed. Di-isopropyl fluorophosphates (DFP, a nerve gas) prevents impulse transfer by combining irreversibly with amino acid serine of acetylcholine esterase.

It also poisons a number of other enzymes like trypsin, chymotrypsin, phosphoglucomutase, elastase, etc. lodoacetamide inhibits enzymes having sulphahydryl (—SH) or imidazole group.

iv. Allosteric Modulation or Feed Back Inhibition:

It is a type of reversible inhibition found in allosteric enzymes. The inhibitor is non-competitive and is usually a low molecular intermediate or product of a metabolic pathway having a chain of reactions involving a number of enzymes. It is, therefore, also called end product or feedback inhibition.

The inhibitor is also called modulator. Modulator is a substance that attaches with an allosteric enzyme at a site other than catalytic one but influences the latter, either inhibiting or activating the same. An example of feed back or allosteric inhibition is stoppage of activity of enzyme hexokinase (glucokinase) by glucose-6-phosphate, the product of reaction catalysed by it (Fig. 9.38).

Another example is inhibition of threonine deaminase by isoleucine (Fig. 9.3). Amino acid isoleucine is formed in bacterium Escherichia coli in a 5-step reaction from threonine. Each step requires a separate enzyme. When isoleucine accumulates beyond a threshold value, its further production stops.

Isoleucine added to the medium of bacterium also stops its internal production showing that its excess prevents some step of the reaction. The latter was found out to be enzyme threonine deaminase which is involved in the first step of the reaction (threonine to a-ketobutyrate).

(i) It has a regulatory role on enzyme activity,

(ii) Enzyme inhibitors have been used in the study of metabolic pathways,

(iii) Some inhibitors are used in controlling pathogenic activity, e.g., sulpha drugs,

(iv) Use of inhibitors have shown the mechanism of enzyme action.

Qualitative feast for carbohydrates/Sugar:

(a) Tests for Glucouse and Fructose:

1. Grape Juice/Fruit Juice contains glucose and fructose. Their presence can be tested by Fehling’s test. Take 5 ml of fruit juice in a test tube. Add an equal quantity of Fehling solution I and II. Boil. A brick red precipitate of cuprous oxide indicates the presence of glucose or fructose in fruit juice. (Cu +2 (Blue) → Cu + red)

2. Fructose gives a red colour with Selivenoff’s reagent (resorcinol + conc. HCI) while glucose gives no colouration.

3. Glucose chars only on heating by the action of conc. H2SO4 while fructose chars in cold.

4. To 5 cc of Benedict’s solution, add 3cc of glucose solution. On boiling, green/red yellow/rust brown precipitate is formed.

5. Sucrose gives negative with Fehling’s test. It is first hydrolysed by conc. H2SO4. Boil 5cc of sucrose solution with few drops of conc. H2SO4. It is then neutralized with NaOH. Boil again and then add Fehling’s solution I and II drop by drop. A brick red precipitate is formed due to hydrolysis of sucrose into glucose and fructose by H2SO4.

6. Stain the pulp of apple/melon with methylene blue. Pectic substances in cell wall are stained violet.

(b) Tests for fats and oils/lipids:

1. Fats are insoluble in water but soluble in ether/acetone/benzene.

2. To few ml of castor oil, add few drops of Sudan III solution. A reddish colour appears.

3. Dissolve fat in alcohol and add a few drops of distilled water. An emulsion of fat in water appears on the surface. Now add few drops of 0.1% sudan III solution made in alcohol. The emulsion turns red.

4. Take 50 ml of fat sample. Add to it 100 ml of 10% NaOH. Boil for 30 minutes. Divide it into two parts A and B. To A part, add few drops of conc. H2SO4. A soapy layer collects at the surface. To В part add saturated solution of NaCI gradually. Soap precipitates and rises to the surface.

5. Test for glycerol: Mix 1 ml of 1% CuS04 solution and 5 drops of the glycerol. To it add 5 drops of 10% NaOH solution. Blue colour is obtained.

(c) Tests for Proteins:

(i) Xanthoproteic test. To 5 cc of protein solution, add 2cc of conc. HNO3. Heat with boiling. A yellow colour appears. Cool it and add excess of 20% NH4OH or NaOH. Orange colour is formed. It is due to nitration of phenolic groups attached to side chains of aromatic amino acids. This test is performed for proteins containing aromatic amino acids like tyrosine, tryptophan etc.

(ii) Grind protein rich seeds (Gram, Pea) with water to make protein solution. Add Millons reagent and heat to boiling. Red colour is formed.

(iii) Biuret Test. To 3 ml of protein solution, add 1 ml of 40% NaOH and few drops of 5% CuSO4 solution. Shake and keep for 15 minutes at room temperature. A violet/pink colour appears which gradually changes to blue and purple. Actually Cu ++ reacts with CONH of proteins and forms a violet coloured complex called biuret (CONH2 — NH—CONH2).

(iv) Beat white of egg with 8 times its volume of H20. Filter and perform xanthoproteic test.

(d) Test for Amino Acids:

Ninhydrin test is best to detect amino acids. Most of the amino acids give purple colour with ninhydrin but tyrosine, phenylalanine aspartic acid give blue colour tryptophan produces olive brown and proline gives yellow colour. 0. 2% ninhydrin solution made in 70% alcohol is put on a piece of filter paper and dried.

Now put a drop of aqueous solution of amino acids on this dry filter paper. Dry in oven. A blue/violet/purple colour appears due to the reaction of ninhydrin with amino group of amino acid to form Ruhemann’s purple compound.

(e) Test for Urine:

Urine is tested for presence of urea, uric acid, creatinine, minerals (Cl, SO4, Ca, PO4), sugar and albumin protein.

(i) Urine for presence of urea: Urea when heated decomposes with the liberation of ammonia and the formation of biuret. The biuret is dissolved in water and develops a violet colour forming a complex with alkaline CuSO4 solution.

Place a small amount of urea crystals in a dry test tube and heat it in a low flame. Urea melts and solidifies. (In case of urine, urine is heated to solidify). Cool the test tube. Add 3 ml of water and shake. Add 1 ml of dil NaOH and 2 drops of 1% CuSO4 solution. Pink colour develops indicating presence of urea.

(ii) For detection of sugar in urine perform Benedicts test. To about 5 ml of Benedicts reagent, add 0.5 ml (8 drops) of urine and boil for 2 minutes. A light green/yellow/brick red precipitate indicates the presence of reducing sugar in urine. The intensity of colour depends upon percentage of sugar in urine.

(iii) For albumin in urine, (a) Fill 3/4th of the test tube by urine after filtering it. Heat the upper 1/3rd of the test tube by a small flame. A turbidity is found on the heated portion of the urine. Add a drop of 33% acetic acid to the urine. Phosphate is dissolved but not the albumin protein, (b) Add a few drops of 30% sulphosalicyclic acid to 2 ml of clear filtered urine. A turbidity indicates the presence of albumin.

Note # 17. Feedback Inhibition of Enzymes:

Feedback inhibition (also called end-product inhibition or allosteric modulation) is one in which the end- product of the reaction acts as inhibitor and inhibits the activity of regulatory enzyme, usually, enzyme of the first step of a biosynthetic pathways. In multi-enzyme system synthesis of a product is completed in a number of steps, each step being catalyzed by a specific enzyme.

In some of such systems, the regulatory enzyme is specifically inhibited by the end-product of the pathway whenever the concentration of the end-product exceeds the cell’s requirement. When the reaction catalyzed by the regulatory enzyme is slowed, all subsequent enzymes act at reduced rates due to the depletion of their substrates.

The rate of production of the pathway’s end-product is thereby brought into balance as per the requirement of the cell. Feedback inhibition is beautifully illustrated by biosynthesis of L-isoleucine from L-threonine (Fig. 10-12).

In this system, the first enzyme, threonine dehydratase, is inhibited by L-isoleucine. No other intermediate in the sequence inhibits threonine dehydratase, nor is any other enzyme of the system inhibited by L-isoleucine.

L-isoleucine binds not to the active site, but to regulatory site on the enzyme molecule this is called allosteric modulation. However, when the concentration of the end-product drops sufficiently, the enzyme reactivates and the end-product is resynthesized.

Note # 18. Specificity of Enzyme:

One characte­ristic that distinguishes an enzyme from all other types of catalysts is its substrate specificity (Fig. 8.19). Enzyme specificity is a result of the uniqueness of the active site of each enzyme.

Enzyme specificity is arbitrari­ly grouped as:

i. Absolute specificity:

Enzymes having absolute specificity will catalyse a particular substrate only and will have no catalytic effect on substrates that are closely related. E.g., Urease will catalyse the hydrolysis of urea but not of methyl urea, thiourea or biuret:

ii. Stereo-chemical specificity:

Most enzymes show a markedly high degree of specificity toward one stereo-isomeric form of the substrate, e.g:

(1) Lactic acid dehydro­genase catalyses the oxidation of the L-lactic acid found in muscle cells but not the D-lactic acid found in certain microorga­nisms.

(2) Fumerase adds water to fumaric acid but not to its cis-isomer-maleic acid.

iii. Group specificity:

Enzymes having group specificity are less selective in that they will act upon structurally similar molecules having the same functional groups, e.g., many of the peptidases.

(1) Pepsin will hydrolyse all peptides having adjacent aromatic amino acids.

(2) Carboxy-peptidase attacks peptides from the carboxyl end of the chain, cleaving the amino acids one at a time.

iv. Linkage specificity:

Enzymes hav­ing linkage specificity are the least specific of all, because they will attack a particular kind of chemical bond, irrespective of the struc­tural features in the vicinity of linkage, e.g., lipases catalyse the hydrolysis of ester link­ages in lipids.

Note # 19. Factors Influencing Enzyme Activities:

An enzyme is active within a narrow range of temperature. The temperature at which an enzyme shows its highest activity is called optimum temperature (Fig. 9.29). It generally corresponds to the body temperature of warm blooded animals, e.g., 37°C in human beings. Enzyme activity de­creases above and below this temperature.

Enzyme becomes inactive below minimum temperature and beyond maximum tem­perature. Low temperature preserves the enzymes in the inactive state. Therefore, it is used in preservation of foods inside cold stor­ages.

Low temperature present inside cold storages prevents spoilage of food by two methods:

(i) Inactivity of enzymes present inside food article and

(ii) Non-activity of microbes because their enzymes also become inactive at low temperature.

High temperature destroys enzymes by causing their denaturation. This occurs at 50°C or so. In between the minimum and maximum temperatures, the reaction velocity doubles for every rise in 10°C (general rule of thumb). A time factor appears beyond optimum tempera­ture. Here there is a rise in velocity for a short time followed by a sharp fall.

As opposed to warm blooded or homoeothermal animals (mammals, birds), there are cold blooded or poikilothermal animals (reptiles, amphibians, fishes, invertebrates) whose body temperature rises or falls with that of environmental temperature.

These animals cannot live in very hot or very cold environment as enzyme functioning will be impaired. Because of this reason, frog seeks moist shady environment during summer and lies in an inactive form (hibernation) in the deeper layers of the soil during winter.

Every enzyme has an optimum pH when it is most effective. A rise or fall in pH reduces enzyme activity by changing the degree of ionisation of its side chains.

A change in pH may also start reverse reaction. Fumarase catalyses fumarate → malate at 6.2 pH and reverse at 7.5 pH. Most of the intracellular enzymes function near neutral pH with the exception of several digestive enzymes which work either in acidic range of pH or alkaline, e.g., 2.0 pH for pepsin, 8.5 for trypsin.

iii. Enzyme Concentration:

The rate of a biochemical reaction rises with the increase in enzyme concentration up to a point called lim­iting or saturation point (Fig. 9.30). Beyond this, increase in enzyme concentration has little effect.

iv. Product Concentration:

If the prod­ucts are allowed to remain in the area of the reaction, the rate of forward reaction will fall. Reverse reaction can also start.

They increase activity of enzymes (e.g., chloride for salivary amylase), function as cofactors (e.g., K + , Mn 2+ ) and convert pro-enzymes to enzyme state. HCl of digestive juice changes pro-enzyme pepsinogen to enzyme pepsin. Pepsin also possesses autocatalytic property as it can also change pepsinogen to pepsin state.

vi. Protein Poisons:

Cyanides, azides, iodoacetate, and salts of heavy metals destroy tertiary structure of enzymes by either combining with cofactor or a group of apoenzyme (—SH group, —COOH).

vii. Radiation Energy:

High energy radiations break hydrogen bonds, ionic bonds, and other weak linkages to destroy enzyme structure.

viii. Substrate Concentration:

Increase in substrate concentration increases the rate of reaction.

The enhanced rate is due to two factors:

(a) Occupation of more and more active sites by the substrate molecules

(b) Higher number of collisions between substrate mol­ecules.

The rise in velocity is quite high in the beginning but it decreases progressively with the increase in substrate concentration. If a graph is plotted for substrate concentration versus reaction velocity, it appears as a hyperbolic curve.

A stage is reached where velocity is maximum. It does not increase further by increasing the substrate concentration. At this stage the enzyme molecules become fully saturated and no active site is left free to bind additional substrate molecules. This satura­tion effect is shown by all enzymes. Because of this Victor Henri (1903) proposed the formation of enzyme-substrate complex as an essential step in enzyme catalysis.

Michaelis Constant (Michaelis Menten Constant, Km). It is a mathematical deriva­tion or constant which indicates the sub­strate concentration at which the chemical reaction catalysed by an enzyme attains half its maximum velocity (Fig. 9.31).

Constant was given forth by Leoner Michaelis and Mand Menten (1913). Km or Michaelis Menten constant generally lies between 10- 1 to 10 -6 M. Km indicates affinity of the enzyme for its substrate.

A high Km indicates low affinity while a low Km shows strong affinity. If an enzyme acts on more than one substrate it shows different Km values for them. Thus enzyme protease acts on large number of proteins. Its Km value will differ from protein to protein.

Allosteric enzymes do not show a typical Michaelis Menten constant or behaviour. The classical hyperbolic curve is replaced by a sigmoid saturation curve.

Note # 20. Biochemical Pathways of Enzymes:

All chemical reactions occurring in the living systems are mediated through organic catalysts called enzymes. Like catalysts, enzymes remain unchanged at the end of the chemical reactions. They can, therefore, be used and re­used. Catalysts (e.g., Platinum) are relatively unselective but enzymes are highly specific.

There is a separate enzyme for every biochemical reaction, i.e., there is no un-catalysed metabolic conversion in living systems. Even an otherwise physical process like dissolution of carbon dioxide in water is catalysed in living systems.

It is because rate of enzyme catalysed reaction is hundreds of times higher than the rate of un-catalysed reaction. In the absence of enzyme, nearly 200 molecules of carbon dioxide dissolve in water per hour to form carbonic acid. In the presence of enzyme carbonic anhydrase, some 600,000 molecules of carbonic acid are formed per second. This is an acceleration of about 10 million times.

During chemical reactions, older chemical bonds are broken and new chemical bonds are established, e.g.,

It is an inorganic chemical reaction. Organic chemical reactions also occur similarly. Rate of chemical or a physical reaction is determined by the amount of product formed per unit time.

Rate of chemical reaction doubles or decreases by half for every 10°C rice or fall. It is called general rule of thumb. As there are thousands of chemical reactions occurring in living cells, thousand of different types of enzymes develop inside the cells. Enzymes are generally globular proteins having one or more clefts over their surface. The clefts function as active sites. Active sites attract substrate molecules or reactants.

An enzyme-substrate complex is formed. Chemical reaction occurs in this stage. It forms products. The products leave the active site which becomes free to attract more substrate molecules. The active site of each type of enzyme is specific for its substrate. This explains the specificity of enzymes. Sucrase will act only on sucrose and no other disaccharide.

Metabolism is of two kinds, catabolism and anabolism. Anabolism includes all the “building up” reactions.

It is also called constructive metabolism since it involves the synthesis of complex substances from simpler ones, e.g., synthesis of organic compounds from CO2 and H2O during photosynthesis, formation of starch from glucose, production of proteins from amino acids, formation of lipids from fatty acids and alcohols. Energy is stored (as potential energy) in the process.

Catabolism (= katabolism) constitutes “breakdown reactions”. It is also known as destructive metabolism because it involves breaking of complex substances into simpler ones. Potential energy present in the complex substances is converted into kinetic energy.

Same of energy is trapped as chemical energy in adenosine triphosphate (ATP). The latter is also called energy currency of living systems. Respiration is an example of catabolism. It releases energy for performing different body activities.

Metabolism, therefore, involves changes in energy and materials. Bioenergetics is the scientific study of energy transformations in the living systems, e.g., organisms, ecosys­tems. Biochemical Pathways are Tightly Regulated. Biochemical reactions do not occur singly.

Neither they are unregulated. Unregulated anabolic and catabolic reactions occurring simultaneously can create chemical chaos as a substance synthesized in anabolism will be immediately broken down in catabolism. There can be excess synthesis or breakdown of a material resulting in unnecessary wastage of energy and materials.

This does not happen. Actually cells have well separated biochemical pathways which are under tight control of regulatory systems which include control of enzyme synthesis, activation and inhibition of enzymes and feedback systems.

Most enzymes have allosteric sites away from active sites. Presence of an activator over the allosteric site makes the active site functional while occurrence of an inhibitor over the allosteric site dysfunctions the active site.

Further, a metabolic or biochemical pathway has a number of steps, each controlled by a separate enzyme. The pathway becomes operational only when the substrate of the first reaction and its active enzyme are available. The product of the first reaction generally becomes the substrate of the second reaction catalysed by a separate enzyme.

The product of the second reaction becomes the substrate of the third reaction and so on till the final product is formed. The cells control the amount of final product as per their requirement by either controlling availability of the first substrate, first enzyme, utilisation of intermediate product or feedback mechanism.

In feedback mechanism or inhibition, the excess final product functions as allosteric inhibitor of the first enzyme, e.g., glucose 6-phosphate for enzyme hexokinase.

The Living State:

Hundred and thousands of metabolites or biomolecules occur in organisms in concen­tration characteristics of each of them, e.g., glucose is 4.5-5.0 mM in blood, hormones in Nano gram/ml.

The living systems maintain this concentration of biomolecules because they are in metabolic flux, always remaining in non-equilibrium steady state where equilibrium is seldom achieved. No work can be carried out in equilibrium state.

Therefore, living systems are regularly receiving an input of energy to prevent reaching an equilibrium and remain always in non-equilibrium steady state. Energy is obtained from metabolism. Metabo­lism and living state are, therefore, complementary and synonymous. There cannot be a living state without metabolism.

Note # 21. Regulation of Enzyme:

Biochemical reaction studies have shown that the pace of a chemical reaction in a biological system is maintained by the activities of the enzymes. Enzymes are rather unstable molecules and are synthesized and degraded simultaneously. Their activities may be regulated either through their synthesis or by modifying the existing enzyme molecules.

The activities of enzyme molecules are regulated by several ways which are the following:

I. Allosteric Regulation:

Allosteric regulation is a fine mechanism of controlling a reaction through the enzyme activity. Some enzymes (called allosteric enzymes), show sigmoidal curve between the substrate concentration and the activity. The activity of these enzymes is modified by several metabolites. The effect of different concentrations of ‘activator’ and ‘inhibitor’ on these enzymes is also sigmoid.

These effector molecules have a structure different from the substrate molecules. In most of the cases, allosteric inhibitors are the end products of the reaction inhibiting the first enzyme in the series.

Thus, this kind of inhibition is called feedback inhibition, end product inhibition or retro-inhibition. The allosteric activators arc normally one of the substrates or cofactors of the enzyme. The effect of the allosteric ‘inhibitor’ or ‘activator’ on the enzyme is reversible.

When they are withdrawn, the enzyme resumes the original activity:

i. Allosteric Inhibition:

Inhibition of threonine deaminase by isoleucine is an example of allosteric inhibition. Threonine deaminase deaminates threonine to α-ketrobutyrate. The final product of the reaction is isoleucine.

Whenever the accumulation of isoleucine occurs, conversion of threonine to α-ketobutyrate and consequently formation of other intermediaries in the biosynthesis of isoleucine is stopped. When isolcucine is used up, threonine deaminase is reactivated and reactions for the biosynthesis of isoleucine start again.

ii. Allosteric Activation:

Activation of glycogen synthetase by glucose-6-phosphate is an example of allosteric activation. Another example of allosteric regulation (of both inhibitory and activating type) is observed during Pasteur effect. Pasteur effect is the inhibition of glycolysis and fermentation by oxygen. The molecular basis of this effect is the allosteric inhibition of enzyme phosphofructokinase by ATP and citrate and its activation by AMP.

Like many others, this kind of regulation is also of adaptive significance. As the level of AMP increases due to increased use of ATP in the cell, glycolysis is increased by the activation of phosphofructokinase with the result of more formation of ATP. When ATP level exceeds normal requirement of the cell, inhibition of glycolysis occurs through the same enzyme phosphofructokinase and ATP synthesis is stopped.

iii. Mechanism of Allosteric Regulation:

Regarding the mechanism of allosteric regulation, it is proposed that allosteric enzymes have two active centers one for the substrate and the other for effector. These two sites lie either on same or on two different subunits. Binding of a effector molecule to one type of subunit changes the structure of the enzyme molecule in such a way that binding of the substrate to the other subunit is affected.

To explain the mechanism, an example of allosteric regulation of aspartate transcarbamylase may be cited. Asparate transcarbamylase contains two types of subunits. These two types of subuntis may be split apart by treatment with mercurials with one type retaining the ability to bind with the substrate, whereas the other to recognize the inhibitor.

When these two species of subunits are together (active enzyme molecule), binding of the inhibitor to one type of subunit changes the structure of other subunits in such a way that the binding of the substrate is inhibited. When the subunits containing binding sites for the inhibitor are removed enzyme is not affected by the inhibitor.

Further, it gives a typical Michaelis-Menten curve with the substrate concentration. Similarly, the binding of activator may change the molecular structure in such a way that the binding of substrate is facilitated.

II. Isozyme Formation:

Another phenomenon that controls cellular metabolism is the formation of isozymes (isoenzymes). Isozymes are different physical forms of the same enzyme performing the same general function at different rates. They differ to some extent in their amino acid composition also, so that they may be separated by electrophoresis. Lactate dehydrogenase is a classic example of isozyme formation. It catalyses the oxidation of lactate to pyruvate with the help of NAD + .

Lactate dehydrogenase enzyme is a tetramer composed of two distinct types (H and M types) of subunits.

Depending upon the relative number of two types of subunits, lactate dehydrogenase forms 5 isozymes as follows:

The molecular weight of the enzyme is 13,500 but when it is treated with urea or guanidine hydrochloride, it dissociates into subunits each having a molecular weight of about 35,000. The regulation of different isozymes is different. LD1 (HHHH) type of lactate dehydrogenase is found in the heart muscles.

This species is most active at low pyruvate concentration and is inhibited by high concentrations of pyruvate. LD5 (MMMM) type of enzyme is found in skeletal muscle cells and it remains active at high pyruvate concentrations.

Another example of isozyme formation is that of aspartokinase. This enzyme catalyzes the reaction between aspartic acid and ATP to form aspartyl phosphate. Amino acids lysine, methionine, and threonine are final products of the reaction.

The enzyme aspartokinase exists in three forms—aspartokinase I, aspartokinase II and aspartokinase III. Aspartokinase I is inhibited by threonine and III by lysine. Aspartokinase II is insensitive to any of these amino acids. Thus, when any one of these amino acids accumulates, the synthesis of the other is affected very little.

III. Multienzyme System:

Some enzymes exist not as individuals but as aggregates of several enzymes and coenzymes. This they do to channel the metabolities in a pathway efficiently. In an aggregate, each component is arranged in a way that the product of one enzyme becomes the substrate for the other and so on.

An example of enzyme aggregation is that of pyruvic acid dehydrogenase of E. coli. This complex consists of three enzymes- pyruvate decarboxylase, dihydrolipoic dehydrogenase and lipoyl reductase transacetylase. The coenzyme associated with the complex are thiamine pyrophosphate (TPP) and flavin adenine dinucleotide (FAD). A schematic diagram of pyruvate dehydrogenase complex is given in Fig. 27.13.

The stepwise reactions catalyzed by this complex may be written as follows:

Pyruvate + thiamine pyrophosphate → α-hydroxythyl thiamine + pyrophosphate + CO2

α -hydroxyethyl thiamine + lipoate → Thiamine pyrophosphate + acetyl dihydrolipoate

Acetyl dihydrolipoate + CoASH → Acetyl CoA + dihydrolipoate

Dihydrolipoate + NAD + /FAD + → Lipoate + NADH/FADH2

IV. Regulation by Adenylate Energy Charge:

The importance of adenosine phosphates in metabolic processes has been well recognised for living systems. The adenylate energy charge is the measure of total pool of adenosine phosphates in the form of ATP, ADP and AMP. D.D. Atkinson (1969) defines adenylate energy charge as follows.

In most systems, an increase in adenylate energy charge in the physiological range results in stimulation of regulatory enzymes. Although it is a well-known phenomenon in animals and microorganisms, some instances have been recorded from plants also.

It has been shown that the adenylate energy charge affects the activity of pyrophosphomevalonate decarboxylase, which is the key enzyme in the biosynthesis of kaurene from mevalonate. An increase in enzyme activity is observed between adenylate energy charge of 0.8 and 1.0.


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