1.2. Energy and Energy Conversions. Chemistry aspects of life organisation

Physicists define energy as the capacity to do work which occurs when a force operates on an object over a distance. In biochemistry, it is more useful to consider energy as the capacity for change. No cell creates energy—all living things must obtain energy from the environment. Indeed, one of the fundamental physical laws is that energy can neither be created nor destroyed. However, energy can be transformed from one type into another and living cells carry out many such energy transformations. Energy transformations are linked to the chemical transformations that occur in cells—the breaking of chemical bonds, the movement of substances across membranes and so forth. Energy changes are related to changes in matter. Energy comes in many forms, such as chemical energy, light energy and mechanical energy. But all forms of energy can be considered as one of two basic types:

Kinetic energy is the energy of movement. This type of energy does work that alters the state or motion of matter. It can exist in the form of heat, light, electric energy and mechanical energy among others.

Potential energy is the energy of state or position—that is, stored energy. It can be stored in chemical bonds, as a concentration gradient and as electric potential among other ways. Water stored behind a dam has potential energy. When the water is released from the dam, some of this potential energy is converted into kinetic energy which can be harnessed to do work.

Likewise, fatty acids store chemical energy in their C—H bonds and C—C bonds and that energy can be released to do biochemical work. Most of living beings consist of cells. In all cells of all organisms, two types of metabolic reactions occur:

Anabolic reactions (anabolism) link together simple molecules to form more complex molecules. The synthesis of a protein from amino acids is an anabolic reaction. Anabolic reactions require an input of energy and capture it in the chemical bonds that are formed.

Catabolic reactions (catabolism) break down complex molecules into simpler ones and release the energy stored in chemical bonds.

Catabolic and anabolic reactions are often linked. The energy released in catabolic reactions is used to drive anabolic reactions—that is, to do biological work. Cellular activities such as growth, movement and active transport of ions across a membrane all require energy and none of them would proceed without a source of energy. In the discussion that follows, will be discovered the physical laws that govern all energy transformations, identified the energy available to do biological work and considered the direction of energy flow. All these processes can be explained by laws of thermodynamics.

The first law: Energy is neither created nor destroyed

Energy can be converted from one form to another. For example, by striking a match, you convert potential chemical energy to light and heat. The first law of thermodynamics states that in any such conversion of energy, it is neither created nor destroyed.

The first law tells us that in any conversion of energy from one form to another, the total energy before and after the conversion is the same. Potential energy in the chemical bonds of carbohydrates and lipids can be converted to potential energy in the form of ATP. This energy can then be used to produce potential energy in the form of concentration gradients established by active transport and can be converted to kinetic energy and used to do mechanical work, such as muscle contraction.

The second law: Not all energy can be used and disorder tends to increase

The second law of thermodynamics states that, although energy cannot be created or destroyed, when energy is converted from one form to another, some of that energy becomes unavailable to do work. In other words, no physical process or chemical reaction is 100 percent efficient and not all the energy released can be converted to work. Some energy is lost to a form associated with disorder. The second law applies to all energy transformations but we will focus here on chemical reactions in living systems.

Not all energy can be used

In any system, the total energy includes the usable energy that can do work and the unusable energy that is lost to disorder: total energy = usable energy + unusable energy.

In biological systems, the total energy is called enthalpy (H). The usable energy that can do work is called free energy (G). Free energy is what cells require for all the chemical reactions of cell growth, cell division and the maintenance of cell health. The unusable energy is represented by entropy (S) which is a measure of the disorder of the system multiplied by the absolute temperature (T). Thus, it is possible to rewrite the word equation above more precisely as H = G + TS,

Or, if rearrange this expression, G = H – TS (usable energy).

Although G, H or S cannot be measured absolutely, it is possible to determine the change in each at a constant temperature. Such energy changes are measured in calories (cal) or joules (J). A change in energy is represented by the Greek letter delta (). For example, the change in free energy (G) of any chemical reaction is equal to the difference in free energy between the products and the reactants, Greaction = Gproducts – Greactants.

If the necessary free energy is not available, the reaction does not occur.

Depending on the sign and magnitude of S, the entire term, T x S, may be negative or positive, large or small. In other words, in living systems at a constant temperature (no change in T), the magnitude and sign of G can depend a lot on changes in entropy. Large changes in entropy make G more negative in value, as it is shown by the negative sign in front of the T x S term. If a chemical reaction increases entropy, its products are more disordered or random than its reactants. If there are more products than reactants, as in the hydrolysis of a protein to its amino acids, the products have considerable freedom to move around. The disorder in a solution of amino acids will be large, compared with that in the protein in which peptide bonds and other forces prevent free movement. Thus, in hydrolysis, the change in entropy (S) will be positive. If there are fewer products and they are more restrained in their movements than the reactants, S will be negative.

For example, a large protein linked by peptide bonds is less free in its movements than a solution of the hundreds or thousands of amino acids from which it was synthesized.

Disorder tends to increase

The second law of thermodynamics also predicts that, as a result of energy conversions, disorder tends to increase. Chemical changes, physical changes and biological processes all tend to increase entropy and therefore tend toward disorder or randomness. This tendency for disorder to increase gives a directionality to physical processes and chemical reactions. It explains why some reactions proceed in one direction rather than another. How does the second law apply to organisms? Consider the human body with its highly complex structures constructed of simpler molecules. This increase in complexity is in apparent disagreement with the second law. But this is not the case. Constructing 1 kg of a human body requires that about 10 kg of biological materials be metabolized and in the process converted to CO₂, H₂O and other simple molecules, and these conversions require a lot of energy. This metabolism creates far more disorder than the order in 1 kg of flesh.

Life requires a constant input of energy to maintain order

There is no disagreement with the second law of thermodynamics. Having seen that the physical laws of energy apply to living things, it will be now considered how these laws apply to biochemical reactions.

Chemical reactions release or take up energy

In cells, anabolic reactions may make a single product, such as a protein (a highly ordered substance), out of many smaller reactants, such as amino acids (less ordered). Such reactions require or consume energy. Catabolic reactions may break down an ordered reactant, such as a glucose molecule, into smaller, more randomly distributed products, such as carbon dioxide and water. Such reactions give off energy. In other words, some reactions release free energy and others take it up. The amount of energy released (–G) or taken up (+G) by a reaction is related directly to the tendency of the reaction to run to completion (the point at which all the reactants are converted to products). Some reactions tend to run toward completion without any input of energy. These reactions, which release free energy, are said to be exergonic and have a negative G. Reactions that proceed toward completion only with the addition of free energy from the environment are endergonic and have a positive G. If a reaction runs exergonically in one direction (from reactant A to product B, for example), then the reverse reaction (B to A) requires a steady supply of energy to drive it. If A→B is exergonic (G < 0), then B →A is endergonic (G > 0). In principle, chemical reactions can run both forward and backward. For example, if a compound A can be converted into a compound B (A→B), then B, in principle, can be converted into A (B →A), although in given concentrations of A and B, only one of these directions will be favored. Think of the overall reaction as resulting from competition between forward and reverse reactions (A~B). Increasing the concentration of the reactants (A) speeds up the forward reaction and increasing the concentration of the products (B) favors the reverse reaction. At some concentration of A and B, the forward and reverse reactions take place at the same rate. At this concentration, no further net change in the system is observable, although individual molecules are still forming and breaking apart. This balance between forward and reverse reactions is known as chemical equilibrium. Chemical equilibrium is a static state, a state of no net change and a state in which G = 0.

There is good reason to believe that life, as it is known to be, cannot exist without water. Animals and plants that live on the Earth’s land masses had to evolve elaborate ways to retain the water that makes up about 70 percent of their bodies. Aquatic organisms living in water do not need these water-retention mechanisms; thus, biologists have concluded that the first living things originated in a watery environment. This environment need not have been the lakes, rivers and oceans with which we are familiar. Living organisms have been found in hot springs at temperatures above the usual boiling point of water, in a lake beneath the frozen Antarctic ice, in water trapped 2 miles below the Earth’s surface, in water 3 miles below the surface of the sea, in extremely acid and extremely salty water, and even in the water that cools the interiors of nuclear reactors. With 20 trillion galaxies in the universe, each with 100 billion stars, there are many planets out there and if our own solar system is typical, some of them have the water needed for life. As biologists contemplate how life could originate from nonliving matter, their attention focuses not just on the presence of water but on what is dissolved in it. A major discovery of biology is that living things are composed of the same types of chemical elements as the vast nonliving portion of the universe. This mechanistic view— that life is chemically based and obeys universal physicochemical laws—is a relatively recent one in human history. The concept of a “vital force” responsible for life, different from the forces found in physics and chemistry, was common in Western culture until the nineteenth century and many people still assume such a force exists. However, most scientists adhere to a mechanistic view of life and it is the cornerstone of medicine and agriculture. Before describing how chemical elements are arranged in living creatures, some fundamental chemical concepts on changes of matter will be considered. In addition to changes in state (solid to liquid to gas), substances undergo changes that transform both their composition and their characteristic properties.

The first chemical signatures indicating the presence of life here are about 4 billion years old. Thus, it took 600 million years, during a geological time frame called the Hadean, for the chemical conditions on Earth to become just right for life. Key among those conditions was the presence of water. The Ancient Earth probably had a lot of water high in the atmosphere. But the new planet was hot and this water evaporated into space. As the Earth cooled, it became possible for water to remain on its surface but where did that water come from? One current view is that comets—loose agglomerations of dust and ice that have orbited the sun since the planets formed—struck the Earth repeatedly and brought not only water but other chemical components of life, such as nitrogen. As the Earth cooled, chemicals from the rocks dissolved in the water and simple chemical reactions took place. Some of these reactions could have led to life but impacts of large comets and rocky meteorites would have released enough energy to heat the developing oceans almost to boiling, thus destroying any early life. These large impacts eventually subsided and life gained a foothold about 3.8 to 4 billion years ago. The prebiotic Hadean was over. The Archean had begun and there has been life on the Earth ever since. Like the rest of the chemical world, living things are made up of atoms and molecules.

1.2.1. Life and Chemistry: Atoms and Small Molecules

Atoms: The Constituents of Matter

More than a trillion (1012) atoms could fit over the period at the end of this sentence. Each atom consists of a dense, positively charged nucleus, around which one or more negatively charged electrons move The nucleus contains one or more protons and may contain one or more neutrons. Atoms and their component particles have volume and mass, which are properties of all matter. Mass measures the quantity of matter present; the greater the mass, the greater the quantity of matter. The mass of a proton serves as a standard unit of measure: the atomic mass unit (amu) or dalton (named after the English chemist John Dalton). A single proton or neutron has a mass of about 1 dalton (Da). Because the mass of an electron is negligible compare to the mass of a proton or a neutron, the contribution of electrons to the mass of an atom can usually be ignored when measurements and calculations are made.

These are electrons, however, that determine how atoms will interact in chemical reactions and they will be discussed extensively later in this chapter. Each proton has a positive electric charge defined as +1 unit of charge. An electron has a negative charge equal and opposite to that of a proton; thus, the charge of an electron is –1 unit. The neutron, as its name suggests, is electrically neutral, so its charge is 0 unit. Unlike charges (+/–) attract each other; like charges (+/+ or –/–) repel each other. Atoms are electrically neutral: the number of protons in an atom equals the number of electrons.

An element is made up of only one kind of atom

An element is a pure substance that contains only one type of atom. The element hydrogen consists only of hydrogen atoms; the element iron consists only of iron atoms. The atoms of each element have certain characteristics or properties that distinguish them from the atoms of other elements. The more than 100 elements, found in the universe, are arranged in the periodic table. These elements are not found in equal amounts. Stars have abundant hydrogen and helium. The earth’s crust and those of the neighboring planets are almost half oxygen, 28 percent silicon, 8 percent aluminum, 2–5 percent each of sodium, magnesium, potassium, calcium and iron, and contain much smaller amounts of the other elements. About 98 percent of the mass of every living organism (bacterium, turnip or human) is composed of just six elements: carbon, hydrogen, nitrogen, oxygen, phosphorus and sulfur. The chemistry of these six elements will be the primary concern here but the others are not unimportant. Sodium and potassium, for example, are essential for nerves to function; calcium can act as a biological signal; iodine is a component of a vital hormone; and plants need magnesium as part of their green pigment (chlorophyll) and molybdenum in order to incorporate nitrogen into biologically useful substances.

The number of protons identifies the element

An element is distinguished from other elements by the number of protons in each of its atoms. This number which does not change, is called the atomic number. An atom of helium has 2 protons, and an atom of oxygen has 8 protons; the atomic numbers of these elements are thus 2 and 8, respectively. Along with a definitive number of protons, every element except hydrogen has one or more neutrons in its nucleus. The mass number of an atom is the total number of protons and neutrons in its nucleus. The nucleus of a helium atom contains 2 protons and 2 neutrons; oxygen has 8 protons and 8 neutrons. Therefore, helium has a mass number of 4 and oxygen a mass number of 16. The mass number may be thought of as the mass of the atom in daltons. Each element has its own one- or two-letter chemical symbol. For example, H stands for hydrogen, He for helium and O for oxygen. Some symbols come from other languages: Fe (from the Latin ferrum) stands for iron, Na (Latin natrium) for sodium and W (German Wolfram) for tungsten.

Electron behavior determines chemical bonding

When considering atoms, biologists are concerned primarily with electrons because the behavior of electrons explains how chemical changes occur in living cells. These changes, called chemical reactions or just reactions, are changes in the atomic composition of substances. The characteristic number of electrons in each atom of an element determines how its atoms will react with other atoms. All chemical reactions involve changes in the relationships of electrons with one another.

The location of a given electron in an atom at any given time is impossible to determine. We can only describe a volume of space within the atom where the electron is likely to be. The region of space where the electron is found at least 90 percent of the time is the electron’s orbital. In an atom, a given orbital can be occupied by at most two electrons. Thus, any atom larger than helium (atomic number 2) must have electrons in two or more orbitals. The different orbitals have characteristic forms and orientations in space. The orbitals, in turn, constitute a series of electron shells or energy levels, around the nucleus. The innermost electron shell consists of only one orbital called an s orbital. Hydrogen (1H) has one electron in its first shell; helium (2He) has two of them. All other elements have two first-shell electrons, as well as electrons in other shells. The second shell is made up of four orbitals (an s orbital and three p orbitals) and hence can hold up to eight electrons. The s orbitals fill with electrons first and their electrons have the lowest energy. Subsequent shells have different numbers of orbitals but the outermost shells usually hold only eight electrons. In any atom, the outermost electron shell determines how the atom combines with other atoms—that is, how the atom behaves chemically. When an outermost shell consisting of four orbitals contains eight electrons, there are no unpaired electrons. Such an atom is stable and will not react with other atoms. Examples of chemically stable elements are helium, neon and argon. Reactive atoms seek to attain the stable condition of having no unpaired electrons in their outermost shells. They attain this stability by sharing electrons with other atoms or by gaining or losing one or more electrons. In either case, the atoms are bonded together. Such bonds create stable associations of atoms called molecules.

A molecule is two or more atoms linked by chemical bonds. The tendency of atoms in stable molecules to have eight electrons in their outermost shells is known as the octet rule. Many atoms in biologically important molecules—for example, carbon (C) and nitrogen (N)—follow the octet rule. However, some biologically important atoms are exceptions to the rule. Hydrogen (H) is the most obvious exception attaining stability when only two electrons occupy its single shell.

1.2.2. Chemical Bonds: Linking Atoms Together

A chemical bond is an attracting force that links two atoms together to form a molecule. There are several kinds of chemical bonds. Firstly covalent bonds will be discussed, the strong bonds that result from the sharing of electrons. Then will be examined other kinds of interactions, including hydrogen bonds that are weaker than covalent bonds but enormously important to biology. Finally, will be considered ionic bonding which is a consequence of the loss or gain of electrons by atoms.

Covalent bonds consist of shared pairs of electrons

When two atoms attain stable electron numbers in their outermost shells by sharing one or more pairs of electrons, a covalent bond forms. Consider two hydrogen atoms in close proximity, each with a single unpaired electron in its outer shell. Each positively charged nucleus attracts the other atom’s unpaired electron but this attraction is balanced by each electron’s attraction to its own nucleus. Thus, two unpaired electrons become shared by both atoms filling the outer shells of both of them. Two atoms are, thus, linked by a covalent bond and a hydrogen gas molecule (H2) is formed. A molecule made up of more than one type of atom is called a compound. A molecular formula uses chemical symbols to identify the different atoms in a compound and subscript numbers to show how many of each type of atoms are present. Thus, the formula for sucrose—table sugar—is C₁₂H₂₂O₁₁. Each compound has a molecular weight (molecular mass) that is the sum of the atomic weights of all atoms in the molecule. Looking at the Periodic table in Annex 3 you can calculate the molecular weight of table sugar to be 342. Molecular weights are usually related to a molecule’s size. A carbon atom has a total of six electrons; two electrons fill its inner shell and four are in its outer shell. Because its outer shell can hold up to eight electrons, carbon can share electrons with up to four other atoms—it can form four covalent bonds. When an atom of carbon reacts with four hydrogen atoms, a molecule called methane (CH₄) forms. Thanks to electron sharing, the outer shell of methane’s carbon atom is filled with eight electrons and the outer shell of each hydrogen atom is also filled. Four covalent bonds—each consisting of a shared pair of electrons—hold methane together. Table shows the covalent bonding capacities of some biologically significant elements.

Orientation of covalent bonds

Covalent bonds are very strong. The thermal energy that biological molecules ordinarily have at body temperature is less than 1 percent of that needed to break covalent bonds. Thus biological molecules, most of which are put together with covalent bonds, are quite stable. This means that their three-dimensional structures and the spaces they occupy are also stable. A second property of covalent bonds is that, for a given pair of atoms, they are the same in length, angle and direction, regardless of the larger molecule of which the particular bond is a part. The four filled orbitals around the carbon nucleus of methane, for example, distribute themselves in space so that the bonded hydrogens are directed to the corners of a regular tetrahedron with carbon in the centre. The three-dimensional structure of carbon and hydrogen is the same in a large, complicated protein as it is in the simple methane molecule. This property of covalent bonds makes the prediction of biological structure possible. Although the orientations of orbitals and the shapes of molecules differ, depending on the types of atoms involved and how they are linked together, it is essential to remember that all molecules occupy space and have three-dimensional shapes. The shapes of molecules contribute to their biological functions.

Unequal sharing of electrons.

If two atoms of the same element are covalently bonded, there is an equal sharing of the pair(s) of electrons in the outer shell. However, when two atoms are of different elements, the sharing is not necessarily equal. One nucleus may exert a greater attractive force on the electron pair than the other nucleus, so that the pair tends to be closer to that atom. The attractive force that an atom exerts on electrons is its electronegativity. It depends on how many positive charges a nucleus has (nuclei with more protons are more positive and thus more attractive to electrons) and how far away the electrons are from the nucleus (closer means more electronegativity).

The closer two atoms are in electronegativity, the more equal their sharing of electrons will be. Looking at the table electronegativities of some elements important in biological systems, it is obvious that two oxygen atoms, both with electronegativity of 3.5, will share electrons equally, producing what is called a nonpolar covalent bond. So will do two hydrogen atoms (both 2.1). But when hydrogen bonds with oxygen to form water, the electrons involved are unequally shared: they tend to be nearer to the oxygen nucleus because it is the most electronegative of two. The result is called a polar covalent bond. Because of this unequal sharing of electrons, the oxygen end of the hydrogen–oxygen bond has a slightly negative charge (symbolized – and spoken as “delta negative,” meaning a partial unit of charge) and the hydrogen end is slightly positive (+). The bond is polar because these opposite charges are separated at the two ends or poles of the bond. The partial charges that result from polar covalent bonds produce polar molecules or polar regions of large molecules. Polar bonds greatly influence the interactions between molecules containing them.

Hydrogen bonds

Hydrogen bonds may form within or between atoms with polar covalent bonds

In liquid water, the negatively charged oxygen atom of one water molecule is attracted to the positively charged hydrogen (+) atoms of another water molecule. (Remember, negative charges attract positive charges.) The bond resulting from this attraction is called a hydrogen bond. Hydrogen bonds are not restricted to water molecules. They may form between an electronegative atom and a hydrogen atom covalently bonded to a different electronegative atom. A hydrogen bond is a weak bond; it has about one-tenth (10%) the strength of a covalent bond between a hydrogen atom and an oxygen atom. However, where many hydrogen bonds form, they have considerable strength and greatly influence the structure and properties of substances. Hydrogen bonds also play important roles in determining and maintaining the three-dimensional shapes of giant molecules such as DNA and proteins.