What is the Function and Scope of Matter and Energy in Biology?

In biology, living systems exhibit a highly organized state of energy and matter. As one may peep into the history of man’s attempts to solve the “mystery of life”, it becomes clear that a direct correlation exists between the concepts formulated and the development of techniques which have made it possible to study life at progressively lower levels of organization.

The whole process is somewhat like opening a whole series of boxes-within-boxes. With the opening of each box, something more is learned about the nature of the entire package-if only the fact that yet another box remains to be opened.

To a great extent, the study of life has been, and continues to be, a matter of opening still smaller boxes. New concepts are formulated at each level of organization, and every new concept becomes fruitful of still other conceptual schemes.

Because modern biology has reached a point at which it is opening exceedingly small boxes, it is necessary that the student of biology know something of the nature of matter and energy. Without these concepts, and evaluations, it is impossible to call modern biology as a rapidly advancing experimental science.

Chemistry and physics have found it useful, for their purposes, to conceive of matter and energy in rather technical terms. Such conceptions must take into account the interconversion of matter and energy, and as a result, it is not always meaningful in these sciences to formulate a precise distinction between them. In living systems, it appears that there is very little inter conversion between matter and energy, and for this reason, we shall consider them as separate and distinct things.

Whether they really are different does not concern us at this point; in fact, such a consideration lies in the realm of theoretical physics, and perhaps ultimately, in philosophy. Thus, any definite object or substance within out material universe which can be apprehended by means of our senses or by instruments comes under this definition, whether living or nonliving.

The term substance is used to describe matter which is uniform throughout, such as sugar, copper, water, and so on. From the negative standpoint, materials such as milk, dirt, air, and wood are not substances, because each is composed of several different kinds of matter.

Matter may exist as a solid, a liquid, or a gas. It is possible to convert most substances from one of these states to either of the others by the addition or subtraction of heat.

Water, for instance, can be made to take the form of ice or steam by this means,. Such a change does not alter the fundamental composition of water; it only alters its physical state. Hence, such a change is called a physical change.

However, if water were subjected to some process by which it could be made to combine with some other substance or to separate into its component parts, such alteration of fundamental composition would be termed a chemical change.

Chemists and physicists, using a variety of substances, have investigated chemical changes exhaustively and have concluded that matter consists of certain fundamental particles called atoms.

For purposes of present definition, an atom may be considered the smallest unit of matter which can enter into chemical changes. Just how large is an atom? Exact methods of computation indicate an almost unbelievable degree of smallness.

It has been estimated that 100 million atoms arranged in a row would measure only an inch. Compared with the number of possible substances which exists, there are relatively few kinds of atoms. To be exact, physical scientists recognize the existence of only ninety- two naturally occurring kinds although others have been produced artificially. Let us suppose that we were able to obtain a substance made up of only one kind of atom.

This would be elementary substance, or as it is generally termed, and element, there being possible only ninety-two such substances in nature. Hence, an element is a substance composed of similar atoms. All these elements have been given names, some of which existed long before the particulate nature of the elements they represent was known.

For purposes of brevity, there are symbols that represent each name. In the main, a symbol represents the first letter or the first and second letters of the English or Latin name of the element. For example, the symbol of the element phosphorus is P, that of calcium Ca, that of copper Cu, that of iron Fe, and so on.

A complete list of the elements and their symbols can be found in any introductory textbook of chemistry. Much evidence indicates that atoms are composed of three primary building blocks: protons, neutrons, and electrons. This is true of all atoms except that of hydrogen, which has no neutron. The protons and neutrons have almost 2,000 times the mass of an electron and are held together very tightly to form the compact nucleus of the atom.

A proton has positive electrical charge and a neutron is neutral, which means that the nucleus has net positive charge. The arrangement of protons and neutrons in the atomic nucleus if completely understood, as is the nature of the energy that binds them together.

An electron has a negative electrical charge. The electrons of an atom move about the positively charged nucleus at varying distances from it, travelling at relatively high velocities. The number of electrons in an atom is ordinarily equal to the number of protons, making the atom neutral with respect to electrical charge.

Variations in the numbers of protons, neutrons, and electrons which compose atoms account for differences in the elements they represent. Primarily, there are three ways of identifying an atom or the element to which it belongs. Probably the simplest and most orderly, at least for reference purposes, is to cite the atomic number. The atomic number of an atom is equal to the number of pro tons in the nucleus.

This means that atomic numbers range from 1 for hydrogen, the simplest atom, to 92 for uranium, the most complex of the naturally occurring atoms. Atoms are also identified by mass numbers, in which case the protons and neutrons are considered to have a mass of one each, and the electrons are considered to have no mass.

For example, carbon-12, which has six protons and six neutrons, has an atomic number of six and a mass number of twelve. As an element, it is frequently represented by the symbol in which case the subscript is the atomic number and the superscript it the mass number. A third means of identifying an atom, and one which is closely related to the concept of atomic mass, is that of atomic weight.

Atomic weights if atoms are relative values determined by comparison of a given element with that of carbon-12. Thus the atomic weight of an element indicates whether it is lighter or heavier than carbon and by how much.

For example, hydrogen is approximately one-twelfth as heavy as carbon-12, and chlorine is almost three times as heavy. As recently as the year 1900, physical scientists assumed that the atoms composing a given element were identical.

Because this assumption could not be reconciled with certain experimental data, however, special attention was given to the matter. It was soon learned that most elements are composed of two or more variant forms of atoms.

The variant forms of given element were named isotopes, and are atoms of the same element with the same atomic numbers but different mass numbers. In other words, they have the same number of protons and electrons, but not the same number of neutrons.

For example, the element chlorine has two naturally occurring isotopic forms. One type of atom has a mass number of 35 and the other type has a mass number of 37. Precise analysis reveals that the proportions of these two isotopes in nature are about 75.4 percent of the lighter atoms.

Consequently, the mass number averages out at 35.453. The atomic weight of a particular element, then, is defined as the average forms of that particular element. The fact that most elements are isotopic is very fortunate for biology.

For example, if an investigator wishes to trace the path of the element carbon in some living systems, he may label some carbon compound with the relatively rare C¹⁴ and determine its pathway or its ultimate fate by these of instruments which are capable of detecting it. Within recent decades, isotopes have become widely used in biological research.

For our purposes, it is useful to view the atom as a miniature solar system in which the nucleus is analogous to the sun and the electrons to its planets. Thus, an atom consists more of space than of anything else. As we mentioned previously, electrons are located at varying distances from the nucleus, about which they travel at high velocities.

We will have more to say about the behaviour of electrons, since they become directly involved in chemical reactions. Before considering this aspect of matter, however, we need to understand some basic concepts regarding energy.

In contrast to matter, energy neither occupies space nor possesses mass. Therefore, it cannot be defined from a material of structural viewpoint; it must instead be defined in operational terms, or in terms of its effect on matter.

Energy is sometimes defined as the capacity to do work. Within this concept, it is useful to classify energy as either potential or kinetic energy. Potential energy is inactive or stored energy. It possesses the capacity to affect matter but it is not in the process of doing so.

In contrast, kinetic energy is energy in action, that is, it is in the process of affecting matter. In the system represented by, a certain amount of potential energy is present. As the boulder rolls down the hill, this potential energy is converted to kinetic energy, and the amount released is approximately equal to the amount originally expended in getting the boulder to the top of the hill. Energy may exist in a number of different forms.

The most common of these are thermal energy, radiant energy, mechanical energy, electrical energy, and chemical energy. In both living and nonliving systems, energy is converted from one type to another and from one form to another.

We have only to consider a very common example of this conversion process to realize that it occurs. In an automobile engine, potential chemical energy is present in the form of gasoline. Upon its ignition by Kinetic electrical energy, it is converted to kinetic thermal energy.

This thermal energy is then partially converted to mechanical energy, which is eventually dissipated as heat, and so on. As we shall see in a later chapter, these same types of conversions and transformations occur in living systems.

Thus, both living and nonliving systems demonstrate the first law of thermodynamics, which states that energy can neither be created nor destroyed but can simply be changed in form. Sometimes this generalization is called the law of conservation of energy.

This concept is of considerable importance to biology, and we shall return to it. One additional concept of energy which is basic to an understanding of chemical reactions in both living and nonliving systems is the second law of thermodynamics.

As a concept, this law accounts for a multitude of complex phenomena, but stated simply, it holds that energy tends to dissipate itself. A good illustration of this may be seen in chemistry, where reactions proceed from high to low energy states.

In other words, the second law relates energy changes in a system to the organization of that system. Placed in this context, it states that there is an increase in entropy – that is, a decrease in organization. Since useful energy is organised energy, an increase in entropy means a decrease in useful energy. How does this concept relate to the study of organisms?

From one viewpoint, life itself might be regarded as a refutation of or an exception to the second law of thermodynamics. If there is a tendency in an isolated system to proceed towards randomness, the implication is that energy must be taken constantly into a living system in order for it to maintain its organization.

As a matter of fact, this is what actually occurs in living systems. The human body, for example, takes in potential chemical energy which ultimately supplies kinetic chemical energy.

These processes enable the body to maintain its organization, that is, it is prevented from wasting away. Thermodynamically, a living system is not qualitatively different from a nonliving system; the difference is a quantitative one, leading to increased complexity in the living system.

Energy conversions in living systems are fantastically numerous and varied, but overall, enough energy is supplied from external sources to defer its progress toward randomness.

When the supply of energy is in sufficient, the organism dies, of course. In brief, although a living system may involve more than just physics and chemistry, we have no reason to believe that it demonstrates a physics and chemistry different from that of involving systems.

Thus organisms exemplify not only the first and second laws of thermodynamics, but all other laws of physics and chemistry as well. As we shall emphasize in a later chapter, the initial energy source for organisms is the sun. Green plants are capable of converting a portion of this radiant energy to kinetic and potential chemical energy. Animals and microorganisms then use these plants as a source of potential energy.

At each step of energy, transfer, there is a considerable loss. The energy lost, including the original radiant energy not utilized by green plants, goes on to a more disorganized state, that is, there is an overall increase in entropy.

The universality of the second law of thermodynamics therefore hinges on whether or not the balance sheet of the entire earth-sun system shows a decrease in free or usable energy. Many physicists feel that this is precisely the case, and they view the universe as analogous to a huge clock which was initially would up and which will eventually unwind it completely. Armed with this brief and very simplified concept of energy, let us return to the nature of atomic structure.

Previously, we depicted the atom as a miniature solar system in which the electrons travel about the nucleus in orbital fashion. In order to understand the conditions which make possible the combination of atoms and the transfer of electrons to and from atoms, it is necessary to recognize that electrons do not revolve about the nucleus in a haphazard fashion.

Rather, there are orbits, or “shells, which are restricted in the number of electrons each can contain. The simplest atom that of hydrogen is characterized by the presence of only one proton in the nucleus and one electron in orbit. The helium atom possesses two protons and two neutrons in the nucleus and two electrons occupying the same shell.

Experimental evidence indicates that the first shell surrounding the nucleus of an atom never contains more than two electrons. In the atom of lithium, for example, which possesses three portions in the nucleus and three electrons in orbit, two of these electrons orbit in the first shell and the third orbits in an outside shell?

This second shell may contain as many as eight electrons. When more than ten electrons are present in the atom, a third shell is established outside the first two.

This third shell may contain as many as eighteen electrons, the fourth shell thirty-two, the fifth shell thirty-two, the sixth shell eighteen, and the seventh shell two. However, not more than eight electrons are contained in whichever is the outermost shell of an atom. The shells formed by orbiting electrons are not so much physical entitles as they are energy levels.

According to this concept, electrons may be viewed as units which possess certain amounts of potential energy, this amount in any particular case being determined by the energy level which the electron occupies in the atom, of we think of the nucleus with its net positive charge as attracting the negatively charged electrons with a certain force, then a theory can be presented in an effort to account for the electron-nucleus relationship: The farther the orbit is from the nucleus, the more potential energy it represents.

Perhaps an analogy will serve to clarify this theory. Imagine a cliff in which successively higher steps are cut, with rocks of equal size being placed in these steps. The higher the rock, the more potential it represents, because it required more kinetic energy to get it there in the first place.

In this analogy, the steps represent different energy levels, or orbital’s, and the rocks represent electrons. The analogy breaks down somewhat when we consider that electrons are in motion, but the principle is the same. However, because they are in motion, those electrons farthest from the nucleus can be removed from the influence of the nucleus more easily than those electrons situated closer to the nucleus.

This is because the attracting force is inversely proportional to the square of the distance of the electron from the nucleus. This concept of energy levels in the atom is basic to an understanding of the interaction of atoms to form molecules.

It is also essential to an understanding of the energy transformations which occur in living systems, where electron shifts from one level to another within atoms are accompanied by a gain or loss in energy.

As we shall see in a later chapter, these mechanisms account for the ability of green plants to “capture” the energy of sunlight, and they enable all organisms to make certain transformations within their cells.