Chapter 2. THE MOLECULES OF LIFE

CHAPTER 2 INTRODUCTION

CORE CONCEPTS

2.1 The atom is the fundamental unit of matter.

2.2 Atoms can combine to form molecules linked by chemical bonds.

2.3 Water is abundant and essential for life.

2.4 Carbon is the backbone of organic molecules.

2.5 Organic molecules include proteins, nucleic acids, carbohydrates, and lipids, each of which is built from simpler units.

2.6 Life likely originated on Earth by a set of chemical reactions that gave rise to the molecules of life.

When biologists speak of diversity, they commonly point to the 2 million or so species named and described to date, or to the 10–100 million living species thought to exist in total. Life’s diversity can also be found at a very different level of observation: in the molecules that make up each and every cell. Life depends critically on many essential functions, including establishing a boundary to separate cells from their surroundings, storing and transmitting genetic information, and harnessing energy from the environment. These functions ultimately depend on the chemical characteristics of the molecules that make up organisms.

In spite of the diversity of molecules and functions, the chemistry of life is based on just a few types of molecule, which in turn are made up of just a few elements. Of the 100 or so chemical elements, only about a dozen are found in more than trace amounts in living organisms. These elements interact with one another in only a limited number of ways. So, the question arises: How is diversity generated from a limited suite of chemicals and interactions? The answer lies in some basic features of chemistry.

2.1 PROPERTIES OF ATOMS

Since antiquity, it has been accepted that the materials of nature are made up of a small number of fundamental substances combined in various ways. Aristotle called these substances elements, and recognized four: earth, air, fire, and water. From the seventeenth century through the end of the nineteenth, elements were defined as pure substances that could not be broken down further by the methods of chemistry. In time, it was recognized that each element contains only one type of atom, the basic unit of matter. By 1850, about 60 elements were known, including such common ones as oxygen, copper, gold, and sodium. Today, 118 elements are known, of which 94 occur naturally and 24 have been created artificially in the laboratory. Elements are often indicated by a chemical symbol, which consists of a one-or two-letter abbreviation of the name of the element. For example, carbon is represented by C, hydrogen by H, and helium by He.

2.1.1 Atoms consist of protons, neutrons, and electrons.

Elements are composed of atoms. The atom contains a dense central nucleus made up of positively charged particles called protons and electrically neutral particles called neutrons. A third type of particle, the negatively charged electron, moves around the nucleus at some distance from it. Carbon, for example, typically has six protons, six neutrons, and six electrons (Fig. 2.1).

Figure 2.1: A carbon atom. Each carbon atom has six protons, six neutrons, and six electrons. The net charge of any atom is neutral because there are as many electrons as protons.

The number of protons, or the atomic number, specifies the atom as a particular element; an atom with one proton is hydrogen, for example, and an atom with six protons is carbon. The atomic number is sometimes indicated as a left subscript to the chemical symbol, for example ₆C. Here, the “6” represents the number of protons in a carbon atom.

Together, the protons and neutrons determine the atomic mass, the mass of the atom. Each proton and neutron, by definition, has a mass of 1, whereas an electron has negligible mass. The number of neutrons in atoms of a particular element can vary, changing its mass. Isotopes are atoms of the same element that have different numbers of neutrons. For example, carbon has three isotopes: About 99% of carbon atoms have six neutrons and six protons, for an atomic mass of 12; about 1% has seven neutrons and six protons, for an atomic mass of 13; and only a very small fraction has eight neutrons and six protons, for an atomic mass of 14. The atomic mass is sometimes indicated as a left superscript to the chemical symbol. For instance, ¹²C is the isotope of carbon with six neutrons and six protons.

Typically, an atom has the same number of protons and electrons. Because a carbon atom has six protons and six electrons, the positive and negative charges cancel each other out and the carbon atom is electrically neutral. Some chemical processes cause an atom to either gain or lose electrons. An atom that has lost an electron is positively charged, and one that has gained an electron is negatively charged. Electrically charged atoms are called ions. The charge of an ion is specified in the right superscript position next to the chemical symbol. Thus, H⁺ indicates a hydrogen ion that has lost an electron and is positively charged.

2.1.2 Electrons occupy regions of space called orbitals.

Electrons move around the nucleus, but not in the simplified way shown in Fig. 2.1. The exact path that an electron takes is unknown, but it is possible to identify a region in space, called an orbital, where an electron is present most of the time. For example, Fig. 2.2a shows the orbital for hydrogen, which is simply a sphere occupied by a single electron. Most of the time, the electron is found within the space defined by the sphere, although its exact location at any instant is unpredictable.

Figure 2.2: Electron orbitals and energy levels (shells) for hydrogen and carbon. The orbital of an electron can be visualized as a cloud of points that are more dense where the electron is more likely to be. The hydrogen atom contains a single orbital, in a single energy level (a and c). The carbon atom has five orbitals, one in the first energy level and four in the second energy level (b and c).

Orbitals have certain properties. The maximum number of electrons in any orbital is two. Most atoms have more than two electrons and so have several orbitals positioned at different distances from the nucleus. These orbitals differ in size and shape. Electrons in orbitals close to the nucleus have less energy than do electrons in orbitals farther away, so electrons fill up orbitals close to the nucleus before occupying those farther away. Several orbitals can exist at a given energy level, or shell. The first shell consists of the spherical orbital shown in Fig. 2.2a.

Fig. 2.2b shows electron orbitals for carbon. Of carbon’s six electrons, two occupy the small spherical orbital representing the lowest energy level. The remaining four are distributed among four possible orbitals at the next highest energy level: One of these four orbitals is a sphere (larger in diameter than that at the lowest energy level) and three are dumbbell-shaped. In carbon, the outermost spherical orbital has two electrons, two of the dumbbell-shaped orbitals have one electron each, and one of the dumbbell-shaped orbitals is empty. Because a full orbital contains two electrons, it would take a total of four additional electrons to completely fill all of the orbitals at this energy level. Therefore, after the first shell, the maximum number of electrons per energy level is eight. Fig. 2.2c shows that the highest energy level, or shell, of carbon, represented by the outermost circle, contains four electrons, and that of hydrogen contains one electron.

2.1.3 Elements have recurring, or periodic, chemical properties.

The chemical elements are often arranged in a tabular form known as the periodic table of the elements, shown in Fig. 2.3 and generally credited to the nineteenth-century Russian chemist Dmitri Mendeleev. The table provides a way to organize all the chemical elements in terms of their chemical properties.

Figure 2.3: The periodic table of the elements. Elements are arranged by increasing number of protons, the atomic number. The elements in any column share similar chemical properties.

In the periodic table, the elements are indicated by their chemical symbols and arranged in order of increasing atomic number. For example, the second row of the periodic table begins with lithium (Li) with 3 protons and ends with neon (Ne) with 10 protons.

For the first three horizontal rows in the periodic table, elements in the same row have the same number and types of orbitals. Across a row, therefore, electrons fill the shell until a full complement of electrons is reached on the right-hand side of the table. Fig. 2.4 shows the filling of the shells for the second row of the periodic table. The elements in a vertical column are called a group or family. Members of a group all have the same number of electrons in their outermost orbital. For example, carbon (C) and lead (Pb) both have four electrons in their outermost orbital. The number of electrons in the outermost orbital determines in large part how elements behave and interact with other elements, as we will see in the next section.

Figure 2.4: Energy levels (shells) of row 2 of the periodic table. The complete complement of electrons in the outer shell of this row of elements is 8.

2.2 MOLECULES AND CHEMICAL BONDS

Atoms can combine with other atoms to form molecules, which are substances made up of two or more atoms. When two atoms form a molecule, the individual atoms interact in what is called a chemical bond, a form of attraction between atoms that holds them together. The ability of atoms to form molecules in part explains why just a few types of element can make many different molecules and perform diverse functions in a cell. There are many different ways in which atoms can interact with one another, and therefore many different types of chemical bond, as we describe in this section.

2.2.1 A covalent bond results when two atoms share electrons.

Figure 2.5: A covalent bond. A covalent bond is formed when two atoms share a pair of electrons in a molecular orbital.

The ability of atoms to combine with other atoms is determined in large part by the electrons farthest from the nucleus—those in the outermost orbitals of an atom. These electrons are called valence electrons, and they are at the highest energy level. In many cases, when atoms combine with other atoms to form a molecule, the atoms share valence electrons with each other. Specifically, when the outermost orbitals of two atoms come into proximity, two atomic orbitals each containing one electron merge into a single orbital containing a full complement of two electrons. The merged orbital is called a molecular orbital, and each shared pair of electrons constitutes a covalent bond that holds the atoms together.

Hydrogen gas (H₂), illustrated in Fig. 2.5, is one of the simplest molecules. Note that the right subscript position is used to indicate the number of atoms in a molecule. Each hydrogen atom has a single electron in a spherical orbital. When the atoms join into a molecule, the two orbitals merge into a single molecular orbital containing two electrons that are shared by the hydrogen atoms. A covalent bond between atoms is denoted by a single line connecting the two chemical symbols for the atoms, as shown in the structural formula in Fig. 2.5.

Molecules tend to be most stable when they share enough electrons to completely occupy the outermost energy level or shell. This simple rule of thumb is known as the octet rule and applies to many, but not all, elements. For example, as shown in Fig. 2.6, one carbon atom (C, with four valence electrons) combines with four hydrogen atoms (H, with one valence electron each) to form CH₄ (methane); nitrogen (N, with five valence electrons) combines with three H atoms to form NH₃ (ammonia); and oxygen (O, with six valence electrons) combines with two H atoms to form H₂O (water). Interestingly, the elements of the same column in the next row behave similarly. This is just one example of the recurring, or periodic, behavior of the elements.

Figure 2.6: Four molecules. Atoms tend to combine in such a way as to complete the complement of electrons in the outer shell.

2.2.2 A polar covalent bond is characterized by unequal sharing of electrons.

In hydrogen gas, the electrons are shared equally by the two hydrogen atoms. In many bonds, however, the electrons are not shared equally by the two atoms. A notable example is provided by the bonds in a water molecule (H₂O), which consists of two hydrogen atoms each covalently bound to a single oxygen atom (Fig. 2.7a).

Figure 2.7: Polar covalent bond. Polar covalent bonds do not share the electrons equally.

In a molecule of water, the electrons are not shared equally between the hydrogen and oxygen atoms; rather, the electrons are more likely to be located near the oxygen atom. Unequal sharing of electrons results from a difference in the ability of the atoms to attract electrons, a property known as electronegativity. Electronegativity tends to increase across a row in the periodic table; as the number of protons across a row increases, electrons are held more tightly to the nucleus. Therefore, oxygen is more electronegative than hydrogen and attracts electrons more than does hydrogen. In a molecule of water, oxygen has a slight negative charge, while the two hydrogen atoms have a slight positive charge (Fig. 2.7b). When electrons are shared unequally between the two atoms, the resulting interaction is described as a polar covalent bond.

2.2.3 A hydrogen bond is an interaction of a hydrogen atom and an electronegative atom.

Because the oxygen and hydrogen atoms have slight charges, water molecules orient themselves to minimize the repulsion of like charges so that positive charges are near negative charges. A hydrogen bond results when a hydrogen atom covalently bound to an electronegative atom (such as oxygen or nitrogen) interacts with an electronegative atom of another molecule. In the case of water, a hydrogen atom covalently bound to an oxygen atom is attracted to and interacts with an oxygen atom of another water molecule. The result of many such interactions is a kind of molecular network stabilized by hydrogen bonds. Typically, a hydrogen bond is depicted by a dotted line, as in Fig. 2.8.

Figure 2.8: Hydrogen bonds in liquid water. Because of thermal motion, hydrogen bonds in water are continually breaking and reforming between different pairs of molecules.

Hydrogen bonds are much weaker than covalent bonds, but it is hydrogen bonding that gives water many of its unusual properties, which are described in the next section. In addition, many weak hydrogen bonds can help stabilize biological molecules, as in the case of nucleic acids and proteins.

2.2.4 An ionic bond forms between oppositely charged ions.

Figure 2.9: An ionic bond.(a) Sodium chloride (salt) is formed by the attraction of two ions. (b) In solution, the ions are surrounded by water molecules.

In water, the difference in electronegativity between the oxygen and the hydrogen atoms leads to unequal sharing of electrons. In more extreme cases, when an atom of very high electronegativity is paired with an atom of very low electronegativity, the difference in electronegativity is so great that the electronegative atom “steals” the electron from its less electronegative partner. In this case, the atom with the extra electron has a negative charge and is a negative ion. The atom that has lost an electron has a positive charge and is a positive ion. The two ions are not covalently bound, but they associate with each other because of the attraction of opposite charges in what is called an ionic bond. An example of a compound formed by the attraction of a positive ion and a negative ion is table salt, or sodium chloride (NaCl) (Fig. 2.9a).

When sodium chloride is placed in water, the salt dissolves to form sodium ions (written as “Na⁺”) that have lost an electron and so are positively charged, and chloride ions (Cl⁻) that have gained an electron and are negatively charged. In solution, the two ions are pulled apart and become surrounded by water molecules: The negatively charged ends of water molecules are attracted to the positively charged sodium ion and the positively charged ends of other water molecules are attracted to the negatively charged chloride ion (Fig. 2.9b). Only as the water evaporates do the concentrations of Na⁺ and Cl⁻ increase to the point where the ions join and precipitate as salt crystals.

2.2.5 A chemical reaction involves breaking and forming chemical bonds.

Figure 2.10: A chemical reaction. During a chemical reaction, atoms retain their identity, but their connections change as bonds are broken and new bonds are formed.

The chemical bonds that link atoms in molecules can change in a chemical reaction, a process by which given molecules, called reactants, are transformed into different molecules, called products. During a chemical reaction, atoms keep their identity but change their chemical bonds.

For example, two molecules of hydrogen gas (2H₂) and one molecule of oxygen (O₂) can react to form two molecules of water (2H₂O), as shown in Fig. 2.10. In this reaction, the numbers of each type of atom are conserved, but their arrangement is different in the reactants and the products. Specifically, the H–H bond in hydrogen gas and the OO bond in oxygen are broken. At the same time, each oxygen atom forms new covalent bonds with two hydrogen atoms, forming two molecules of water. In fact, this reaction is the origin of the name “hydrogen,” which literally means “water former.” The reaction releases a good deal of energy; it was used in the main engine of the Space Shuttle.

In biological systems, chemical reactions provide a way to build and break down molecules for use by the cell, as well as to harness energy, which can be stored in chemical bonds (Chapter 6).

2.3 WATER: THE MEDIUM OF LIFE

On Earth, all life depends on water. Indeed, life originated in water, and the availability of water strongly influences the environmental distributions of different species. Furthermore, water is the single most abundant molecule in all cells, so water is the medium in which the molecules of life interact. In the late 1990s, the National Aeronautic and Space Administration (NASA) announced that the search for extraterrestrial life would guide continuing exploration of the solar system and beyond. NASA’s operational strategy was simple: Follow the water. NASA’s logic was straightforward: Within our solar system, Earth stands out both for its abundance of water and the life it supports. What makes water so special as the medium of life?

2.3.1 Water is a polar molecule.

As we saw earlier, water molecules have polar covalent bonds, characterized by an uneven distribution of electrons. A molecule like water that has regions of positive and negative charge is called a polar molecule. Molecules, or even different regions of the same molecule, fall into two general classes, depending on how they interact with water: hydrophilic (“water loving”) and hydrophobic (“water fearing”).

Hydrophilic compounds, like water itself, are polar; they dissolve readily in water. That is, water is a good solvent, capable of dissolving many substances. Think of what happens when you stir a teaspoon of sugar into water: The sugar seems to disappear as the sugar dissolves. What is happening is that the sugar molecules are dispersing through the water and becoming separated from one another. Sugar is in solution in the watery, or aqueous, environment.

By contrast, hydrophobic compounds are nonpolar and arrange themselves to minimize their contact with water. For example, when oil and water are mixed, oil molecules organize themselves into droplets that limit the oil–water interface. This hydrophobic effect, in which polar molecules like water exclude nonpolar ones, drives such biological processes as the formation of cell membranes (Chapter 3) and the folding of proteins (Chapter 5).

2.3.2 pH is a measure of the concentration of protons in solution.

A small proportion of the molecules in water exist as protons (H⁺) and hydroxide ions (OH⁻). The pH of a solution measures the proton concentration ([H⁺]) by the following formula:

pH = −log [H⁺]

The pH of a solution can range from 0 to 14. Since the pH scale is logarithmic, a difference of one pH unit corresponds to a tenfold difference in hydrogen ion concentration. A solution is neutral (pH = 7) when the concentrations of protons (H⁺) and hydroxide ions (OH⁻) are equal. When the concentration of protons is higher than that of hydroxide ions, then the pH is lower than 7 and the solution is acidic. When the concentration of protons is lower than that of hydroxide ions, then the pH is higher than 7, and the solution is basic. An acid therefore is a molecule that releases a proton (H⁺), and a base is a molecule that accepts a proton in aqueous solution.

Pure water has a pH of 7—that is, it is neutral, with an equal concentration of protons and hydroxide ions. The pH of most cells is approximately 7 and is tightly regulated, as most chemical reactions can be carried out only in a narrow pH range. Certain cellular compartments, however, have a much lower pH. The pH of blood is slightly basic, with a pH around 7.4. This value is sometimes referred to in medicine as physiological pH, as it can change in response to certain diseases. Freshwater lakes, ponds, and rivers tend to be slightly acidic because of dissolved carbon dioxide from the air, which forms carbonic acid in water.

2.3.3 Hydrogen bonds give water many unusual properties.

Water is also characterized by extensive hydrogen bonding, as we have seen. Hydrogen bonds influence the structure of both liquid water (Fig. 2.11a) and ice (Fig. 2.11b). When water freezes, most water molecules become hydrogen bonded to four other water molecules, forming an open lattice-like, crystalline structure we call ice. As the temperature increases and the ice melts, some of the hydrogen bonds are destabilized. This allows the water molecules to pack more closely, and it is the reason why liquid water is more dense than solid water. As a result, ice floats on water, and ponds and lakes freeze from the top down, and therefore do not freeze completely. This special property allows fish and aquatic plants to survive winter in the cold water under the layer of ice.

Figure 2.11: Liquid water and ice. Hydrogen bonds create a dense structure in water (a), and a highly ordered, less dense, crystalline structure in ice (b).

Hydrogen bonds also make water molecules cohesive, meaning that they tend to stick to one another. A consequence of cohesion is high surface tension, a measure of the difficulty of breaking the surface of a liquid. Water cohesion and surface tension contribute to water movement in plants. As water evaporates from leaves, water is pulled upward, sometimes as high as 300 feet above the ground in giant sequoia and coast redwood trees, which are among the tallest trees on Earth.

The hydrogen bonds of water also influence how water responds to heating. Molecules are in constant motion, and this motion increases as the temperature increases. When water is heated, however, the increased motion first breaks hydrogen bonds, and only afterward leads to a temperature increase. The need to break hydrogen bonds first means that water resists temperature changes more than do other substances, a property that is important for living organisms on a variety of scales. In the cell, water resists temperature variations that would otherwise result from numerous biochemical reactions. On a global scale, the oceans minimize temperature fluctuations, stabilizing the temperature on Earth in a range compatible with life.

In short, water is clearly the medium of life on Earth, but is this because water is uniquely suited for life, or is it because life on Earth has adapted through time to a watery environment? We don’t know the answer, but probably both explanations are partly true. Chemists have proposed that under conditions of high pressure and temperature, other small molecules, among them ammonia (NH₃) and some simple carbon-containing molecules, might display similar characteristics friendly to life. However, under the conditions of pressure and temperature that exist on Earth, water is the only molecule uniquely suited to life. Water is a truly remarkable substance, and life on Earth would not be possible without it. Leonardo da Vinci once wrote that water is the driving force of all nature. Without a doubt, water is the driving force of all biology.

2.4 CARBON: LIFE’S CHEMICAL BACKBONE

Hydrogen and helium are far and away the most abundant elements in the universe. In contrast, the solid Earth is dominated by silicon, oxygen, aluminum, iron, and calcium (Chapter 1). In other words, Earth is not a typical sample of the universe. Nor, as it turns out, is the cell a typical sample of the solid Earth. Fig. 2.12 shows the relative abundance by mass of chemical elements present in human cells after all the water has been removed. Note that just four elements—carbon (C), oxygen (O), hydrogen (H), and nitrogen (N)—constitute 94% of the total dry mass, and that the most abundant element is carbon. The elemental composition of human cells is typical of all cells. Human life, and all life as we know it, is based on carbon. Carbon molecules play such an important role in living organisms that carbon-containing molecules have a special name—they are called organic molecules. Their central role in life implies that there must be something very special about carbon, and there is. Carbon has the ability to combine with many other elements to form a wide variety of molecules, each specialized for the functions it carries out in the cell.

Figure 2.12: Approximate proportions by dry mass of chemical elements found in human cells.

2.4.1 Carbon atoms form four covalent bonds.

Figure 2.13: A carbon atom with four covalent bonds. One carbon atom combines with four hydrogen atoms to form methane.

One of the special properties of carbon is that, in forming molecular orbitals, a carbon atom behaves as if it had four unpaired electrons. This behavior occurs because one of the electrons in the outermost sphere moves into the empty dumbbell-shaped orbital. In this process, the single large spherical orbital and three dumbbell-shaped orbitals change shape, becoming four equivalent hybrid orbitals each with one electron.

Fig. 2.13 shows the molecular orbitals that result when one atom of carbon combines with four atoms of hydrogen to form the gas methane (CH₄). Each of the four valence electrons of carbon shares a new molecular orbital with the electron of one of the hydrogen atoms. These bonds can rotate freely about their axis. Furthermore, because of the shape of the orbitals, the carbon atom lies at the center of a three-dimensional structure called a tetrahedron, and the four molecular orbitals point toward the four corners of this structure. The ability of carbon to form four covalent bonds, the spatial orientation of these bonds in the form of a tetrahedron, and the ability of each bond to rotate freely all contribute importantly to the structural diversity of carbon-based molecules.

2.4.2 Carbon-based molecules are structurally and functionally diverse.

Figure 2.14: Diverse carbon-containing molecules. (a) The molecule ethane contains one C–C bond. (b) A linear chain of carbon atoms and a ring structure that also contains oxygen contain multiple C—C bonds.

Carbon has other properties that contribute to its ability to form a diversity of molecules. For example, carbon atoms can link with each other by covalent bonds to form long chains. These chains can be branched, or two carbons at the ends of the chain or within the chain can link to form a ring structure.

Among the simplest chains is ethane, shown in Fig. 2.14a, in which two carbon atoms become connected by a covalent bond. In this case, the orbitals of unpaired electrons in two carbon atoms merge to create the covalent bond. Some more complex examples of carbon-containing molecules, and an abbreviated form, are shown in Fig. 2.14b.

Two adjacent carbon atoms can also share two pairs of electrons, forming a double bond, as shown in Fig. 2.15. Note that each carbon atom has exactly four covalent bonds, but in this case two are shared between adjacent carbon atoms. The double bond is shorter than a single bond and is not free to rotate so that all of the covalent bonds in the carbon atoms connected by a double bond are in the same geometrical plane. As with single bonds, double bonds can be found in chains of atoms or cyclical structures. A double bond between atoms is represented by a double line connecting the two chemical symbols for the atoms.

While the types of atoms making up a molecule help characterize the molecule, the spatial arrangement of atoms is also important. For example, 6 carbon atoms, 13 hydrogen atoms, 2 oxygen atoms, and 1 nitrogen atom can join covalently in many different arrangements to produce molecules with different structures. Two of these many arrangements are shown in Fig. 2.16. Note that some of the connections between atoms are identical (black) and some are different (green), even though the chemical formulas are the same (C₆H₁₃O₂N₁). Molecules that have the same chemical formula but different structures are known as isomers.

Figure 2.15: Molecules containing double bonds between carbon atoms.

Figure 2.16: The isomers isoleucine and leucine. Isoleucine and leucine are isomers: Their chemical formulas are the same, but their structures differ.

We have seen that carbon-containing molecules can adopt a wide range of arrangements, a versatility that helps us to understand how a limited number of elements can create an astonishing variety of molecules. We might ask whether carbon is uniquely versatile. Put another way, if we ever discover life on a distant planet, will it be carbon based, like us? Silicon, just below carbon in the periodic table (see Fig. 2.3), is the one other element that is both reasonably abundant and characterized by four atomic orbitals with one electron each. Some scientists have speculated that silicon might provide an alternative to carbon as a chemical basis for life. However, silicon readily binds oxygen. On Earth, nearly all of the silicon atoms found in molecules are covalently bound to oxygen. Studies of Mars and of meteorites show that silicon is tightly bound to oxygen throughout our solar system, and that is likely to be true everywhere we might explore. There are about 1000 different silicate minerals on Earth, but this diversity pales before the millions of known carbon-based molecules. If we ever discover life beyond Earth, very likely its chemistry will be based on carbon.

2.5 ORGANIC MOLECULES

The chemistry of carbon creates diverse geometries, which in turn lead to molecules with different structures and functions. These molecules perform essential tasks of the cell, including the establishment of a boundary to separate the inside of the cell from the environment, the storage and transmission of genetic information, and the capture, storage, and utilization of energy from the environment. These processes depend on just a few classes of carbon-based molecules. Proteins provide structural support and act as catalysts that facilitate chemical reactions. Nucleic acids encode and transmit genetic information. Carbohydrates provide a source of energy and make up the cell wall in bacteria, plants, and algae. Lipids make up cell membranes, store energy, and act as signaling molecules.

These molecules are all relatively large, and most are polymers, complex molecules made up of repeated simpler units connected by covalent bonds. Proteins are polymers of amino acids, nucleic acids are made up of nucleotides, and carbohydrates such as starch are built from simple sugars. Lipids are a bit different, as we will see—they are defined by a property rather than by their chemical structure. The lipid membranes that define cell boundaries consist of fatty acids bonded to other organic molecules.

Building macromolecules from simple, repeating units provides a means of generating virtually limitless chemical diversity. Indeed, in macromolecules, the building blocks of polymers play a role much like that of the letters in words. In written language, a change in the content or order of letters changes the meaning of the word (or renders it meaningless). For example, by reordering the letters of the word “SILENT” you can write “LISTEN,” a word with a different meaning. Similarly, rearranging the building blocks that make up macromolecules provides an important way to make a large number of diverse macromolecules.

In the following section, we focus on the building blocks of these four key molecules of life, reserving a discussion of the structure and function of the macromolecules for later chapters.

2.5.1 Proteins are composed of amino acids.

Proteins do much of the cell’s work. Some function as catalysts that accelerate the rates of chemical reactions (in which case they are called enzymes), and some act as structural components necessary for cell shape and movement. Your body contains many thousands of proteins that perform a wide range of functions. Since proteins consist of amino acids linked covalently to form a chain, we need to examine the chemical features of amino acids to understand the diversity and versatility of proteins.

Figure 2.17: Amino acids and peptide bonds (a) An amino acid contains four groups attached to a central carbon. (b) Peptide bonds link amino acids to form a protein.

The general structure of an amino acid is shown in Fig. 2.17a. Each amino acid contains a central carbon atom, called the α (alpha) carbon, covalently linked to four groups: a carboxyl group (COOH; red), an amino group (NH₂; blue), a hydrogen atom (H), and an R group, or side chain, (green) that differs from one amino acid to the next. The identity of each amino acid is determined by the structure and composition of the side chain. The side chain of the amino acid glycine is simply H, for example, and that of alanine is CH₃. In most amino acids, the α-carbon is covalently linked to four different groups. Glycine is the exception, since its R group is a hydrogen atom.

Amino acids are linked in a chain to form a protein (Fig. 2.17b). The carbon atom in the carboxyl group of one amino acid is joined to the nitrogen atom in the amino group of the next by a covalent linkage called a peptide bond. In Fig. 2.17b, the chain of amino acids includes four amino acids, and the peptide bonds are indicated in red. The formation of a peptide bond involves the loss of a water molecule since in order to form a C–N bond, the carbon atom must release an oxygen atom and the nitrogen must release two hydrogen atoms. These can then combine to form a water molecule (H₂O). The loss of a water molecule is a consistent feature in the linking of subunits to form polymers such as nucleic acids and complex carbohydrates.

Cellular proteins are composed of 20 amino acids, which can be classified according to the chemical properties of their side chains. The particular sequence, or order, in which amino acids are present in a protein determines how it folds into its three-dimensional structure. The three-dimensional structure, in turn, determines the protein’s function. In Chapter 4, we will examine how the sequence of amino acids in a particular protein is specified and discuss how proteins fold into their three-dimensional configuration.

2.5.2 Nucleic acids encode genetic information in their nucleotide sequence.

Figure 2.18: A ribonucleotide and a deoxyribonucleotide, the units of RNA and DNA.

Nucleic acids are examples of informational molecules—that is, they are large molecules that carry information in the sequence of nucleotides that make them up. This molecular information is much like the information carried by the letters in an alphabet, but in the case of nucleic acids, the information is in chemical form.

The nucleic acid deoxyribonucleic acid (DNA) is the genetic material in all organisms. It is transmitted from parents to offspring, and it contains the information needed to specify the amino acid sequence of all the proteins synthesized in an organism. The nucleic acid ribonucleic acid (RNA) has multiple functions; it is a key player in protein synthesis and the regulation of gene expression.

DNA and RNA are long molecules consisting of nucleotides bonded covalently one to the next. Nucleotides, in turn, are composed of three components: a 5-carbon sugar, a nitrogen-containing compound called a base, and one or more phosphate groups (Fig. 2.18). The sugar in RNA is ribose, and the sugar in DNA is deoxyribose. The sugars differ in that ribose has a hydroxyl (OH) group on the second carbon (designated the 2′ carbon), while deoxyribose has a hydrogen atom at this position (hence, deoxyribose). (By convention, the carbons in the sugar are numbered with primes—1′, 2′, etc.—to distinguish them from carbons in the base—1, 2, etc.)

The bases are built from nitrogen-containing rings and are of two types. The pyrimidine bases (Fig. 2.19a) have a single ring and include thymine (T), cytosine (C), and uracil (U). The purine bases (Fig. 2.19b) have a double-ring structure and include adenine (A) and guanine (G). DNA contains the bases A, T, G, and C, and RNA contains the bases A, U, G, and C. Just as the order of amino acids provides the unique information carried in proteins, so, too, does the sequence of nucleotides determine the information in DNA and RNA molecules.

Figure 2.19: Pyrimidine bases (a) and purine bases (b). Pyrimidines have a single-ring structure, and purines have a double-ring structure.

In DNA and RNA, each adjacent pair of nucleotides is connected by a phosphodiester bond, which forms when a phosphate group in one nucleotide is covalently joined to the sugar unit in another nucleotide (Fig. 2.20). As in the formation of a peptide bond, the formation of a phosphodiester bond involves the loss of a water molecule.

Figure 2.20: The phosphodiester bond. Phosphodiester bonds link successive deoxyribonucleotides, forming the backbone of the DNA strand.

DNA in cells usually consists of two strands of nucleotides twisted around each other in the form of a double helix (Fig. 2.21a). The sugar–phosphate backbones of the strands wrap like a ribbon around the outside of the double helix, and the bases are pointed inward. The bases form specific purine–pyrimidine pairs that are said to be complementary: Where one strand carries an A, the other carries a T; and where one strand carries a G, the other carries a C (Fig. 2.21b). Base pairing results from hydrogen bonding between the bases (Fig. 2.21c).

Figure 2.21: The structure of DNA. (a) DNA is most commonly in the form of a double helix, with the sugar and phosphate groups forming the backbone and the bases oriented inward. (b) The bases are complementary, with A pairing with T and G pairing with C. (c) Base pairing results from hydrogen bonds.

Genetic information in DNA is contained in the sequence, or order, in which successive nucleotides occur along the molecule. Successive nucleotides along a DNA strand can occur in any order, and hence a long molecule could contain any of an immense number of possible nucleotide sequences. This is one reason why DNA is an efficient carrier of genetic information. In Chapter 3, we consider the structure and function of DNA and RNA in greater detail.

2.5.3 Complex carbohydrates are made up of simple sugars.

Figure 2.22: Structural formulas for some 6-carbon aldoses and ketoses.

Many of us, when we feel tired, reach for a candy bar for a quick energy boost. The quick energy in a candy bar comes from sugars, which are quickly broken down to release energy. Sugars belong to a class of molecules called carbohydrates, distinctive molecules composed of C, H, and O atoms, usually in the ratio 1:2:1. Carbohydrates provide a principal source of energy for metabolism.

The simplest carbohydrates are sugars (also called saccharides). Simple sugars are linear or, far more commonly, cyclic molecules containing five or six carbon atoms. All six-carbon sugars have the same chemical formula (C₆H₁₂O₆) and differ only in configuration. Glucose (the product of photosynthesis), galactose (found in dairy products), and fructose (a commercial sweetener) are examples; they share the same formula (C₆H₁₂O₆) but differ in the arrangement of their atoms (Fig. 2.22).

A simple sugar is also called a monosaccharide (mono means “one”), and two simple sugars linked together by a covalent bond is called a disaccharide (di means “two”). Sucrose (C₁₂H₂₂O₁₁), or table sugar, is a disaccharide that combines one molecule each of glucose and fructose. As noted above, nucleotides contain a 5-carbon sugar, either ribose (in RNA) or deoyxribose (in DNA). Simple sugars combine in many ways to form polymers called polysaccharides (poly means “many”) that provide long-term energy storage (starch and glycogen) or structural support (cellulose in plant cell walls). Long, branched chains of monosaccharides are called complex carbohydrates.

Let’s take a closer look at monosaccharides, the simplest sugars. Monosaccharides are unbranched carbon chains with either an aldehyde (HCO) or a ketone (CO) group (Fig. 2.22). Monosaccharides with an aldehyde group are called aldoses and those with a ketone group are known as ketoses. In both types of monosaccharides, the other carbons each carry one hydroxyl (OH) group and one hydrogen (H) atom. When the linear structure of a monosaccharide is written with the aldehyde or ketone group at the top, the carbons are numbered from top to bottom.

Figure 2.23: Formation of the cyclic form of glucose.

Virtually all of the monosaccharides in cells are in ring form (Fig. 2.23), not linear structures. To form a ring, the carbon in the aldehyde or ketone group forms a covalent bond with the oxygen of a hydroxyl group carried by another carbon in the same molecule. For example, cyclic glucose is formed when the oxygen atom of the hydroxyl group on carbon 5 forms a covalent bond with carbon 1, which is part of an aldehyde group. The cyclic structure is approximately flat, and you can visualize it perpendicular to the plane of the paper with the covalent bonds indicated by the thick lines in the foreground. The groups attached to any carbon therefore project either above or below the ring. When the ring is formed, the aldehyde oxygen becomes a hydroxyl group.

Monosaccharides, especially 6-carbon sugars, are the building blocks of complex carbohydrates. Monosaccharides are attached to each other by covalent bonds called glycosidic bonds (Fig. 2.24). As with peptide bonds, the formation of glycosidic bonds involves the loss of a water molecule. A glycosidic bond is formed between carbon 1 of one monosaccharide and a hydroxyl group carried by a carbon atom in a different monosaccharide molecule.

Figure 2.24: Glycosidic bonds. Glycosidic bonds link glucose monomers together to form the polysaccharide starch.

Carbohydrate diversity stems in part from the monosaccharides that make them up, similar to the way that protein and nucleic acid diversity stems from the sequence of their subunits. Some complex carbohydrates are composed of a single type of monosaccharide, while others are a mix of different kinds of monosaccharide. Starch, for example, is a sugar storage molecule in plants composed completely of glucose molecules, whereas pectin, a component of the cell wall, contains up to five different monosaccharides.

2.5.4 Lipids are hydrophobic molecules.

Proteins, nucleic acids, and carbohydrates all are polymers made up of smaller, repeating units with a defined structure. Lipids are different. Instead of being defined by a chemical structure, they share a particular property: Lipids are all hydrophobic. Because they share a property and not a structure, lipids are a chemically diverse group of molecules. They include familiar fats that make up part of our diet, components of cell membranes, and signaling molecules. Let’s briefly consider each in turn.

Figure 2.25: Triacylglycerol and its components.

Triacylglycerol is an example of a lipid that is used for energy storage. It is the major component of animal fat and vegetable oil. A triacylglycerol molecule is made up of three fatty acids joined to glycerol (Fig. 2.25). A fatty acid is a long chain of carbons attached to a carboxyl group (COOH) at one end (Figs. 2.25a and 2.25b). Glycerol is a 3-carbon molecule with OH groups attached to each carbon (Fig. 2.25c).

Fatty acids differ in the length (that is, in the number of carbons) of their hydrocarbon chain. Most fatty acids in cells contain an even number of carbons because they are synthesized by the stepwise addition of 2-carbon units. Some fatty acids have one or more carbon–carbon double bonds; these double bonds can differ in number and location. Fatty acids that do not contain double bonds are described as saturated. Because there are no double bonds, the maximum number of hydrogen atoms is attached to each carbon atom, so all of the carbons are said to be “saturated” with hydrogen atoms (Fig. 2.25a). Fatty acids that contain carbon–carbon double bonds are unsaturated (Fig. 2.25b). The chains of saturated fatty acids are straight, while the chains of unsaturated fatty acids have a kink at each double bond.

Figure 2.26: Van der Waals forces. Transient asymmetry in the distribution of electrons along fatty acid chains leads to asymmetry in neighboring molecules, resulting in weak electrostatic attractions.

Triacylglycerols can contain different types of fatty acids attached to the glycerol backbone. They are all extremely hydrophobic and, therefore, triacylglycerols form oil droplets inside the cell. Triacylglycerols are an efficient form of energy storage because by excluding water molecules a large number can be packed into a small volume.

The hydrocarbon chains of fatty acids do not contain polar covalent bonds like those in a water molecule. Their electrons are distributed uniformly over the whole molecule and thus these molecules are uncharged. However, the constant motion of electrons leads to regions of slight positive and negative charges (Fig. 2.26). These charges in turn attract or repel electrons in neighboring molecules, setting up areas of positive and negative charge in those molecules as well. The temporarily polarized molecules weakly bind to one another because of the attraction of opposite charges. These interactions are known as van der Waals forces. The van der Waals forces are weaker than hydrogen bonds, but many of them acting together help to stabilize molecules.

Because of van der Waals forces, the melting points of fatty acids depend on their length and level of saturation. As the length of the hydrocarbon chains increases, the number of van der Waals bonds between the chains also increases, and this in turn increases the melting temperature. Kinks introduced by double bonds reduce the tightness of the molecular packing and lead to fewer intermolecular interactions and a lower melting temperature. Therefore, an unsaturated fatty acid has a lower melting point than a saturated fatty acid of the same length. Animal fats such as butter are composed of triacylglycerols with saturated fatty acids and are solid at room temperature, whereas plant fats and fish oils are composed of triacylglycerols with unsaturated fatty acids and are liquid at room temperature.

Steroids such as cholesterol are a second type of lipid (Fig. 2.27). Like other steroids, cholesterol has a core composed of 20 carbon atoms bonded to form four fused rings, and it is hydrophobic. Cholesterol is a component of animal cell membranes (Chapter 5) and serves as a precursor for the synthesis of steroid hormones such as estrogen and testerone (Chapter 38).

Figure 2.27: The chemical structure of cholesterol.

Phospholipids are a third type of lipid and a major component of the cell membrane. Whereas triacylglycerol is made up of glycerol attached to three fatty acids, most phospholipids are made up of glycerol attached to two fatty acids and a third molecule that contains a phosphate group (Fig. 2.28a). The phosphate “head” group is hydrophilic, while the fatty acid “tails” are hydrophobic. As a result, phospholipids have hydrophobic and hydrophilic groups in the same molecule, giving them an interesting property when placed in water: They form a variety of structures all of which limit the exposure of the hydrophobic tails to water. One important structure is a bilayer, a two-layered structure with the hydrophilic heads pointing outward toward the aqueous environment and the hydrophobic tails oriented inward, away from water (Fig. 2.28b). The formation of lipid bilayers is discussed more fully in the next chapter.

Figure 2.28: Phospholipids.(a) Most phospholipids are made up of glycerol attached to two fatty acids and a phosphate-containing head group. (b) Phospholipids form a bilayer in aqueous solution.

2.6 HOW DID THE MOLECULES OF LIFE FORM?

In Chapter 1, we considered the similarities and differences between living and non-living things. Four billion years ago, however, the differences may not have been so pronounced. Scientists believe that life originated early in our planet’s history by a set of chemical processes that, through time, produced organisms that could be distinguished from their non-living surroundings. How can we think scientifically about one of biology’s deepest and most difficult problems? We can’t reconstruct key events from the geological record—the study of fossils tells us that life already existed when Earth’s oldest surviving fossil-containing rocks were deposited. The alternative is to approach life’s origins experimentally, asking whether chemical reactions likely to have taken place on the early Earth can generate the molecules of life. It is important to note that even the simplest living organisms living today are far more complicated than our earliest ancestors. No one suggests that cells as we know them emerged directly from primordial chemical reactions. Rather, the quest is to discover simple molecular systems able to replicate themselves, while subject to natural selection.

A key starting point is the observation, introduced earlier in this chapter, that the principal macromolecules found in organisms are themselves made of simpler molecules joined together. Thus, if we want to understand how proteins might have emerged on the early Earth, we should begin with the synthesis of amino acids, and if we are interested in nucleic acids, we should focus on nucleotides.

2.6.1 The building blocks of life can be generated in the laboratory.

Research into the origins of life was catapulted into the experimental age in 1953, with an elegant experiment carried out by Stanley Miller, then a graduate student in the laboratory of Nobel laureate Harold Urey. Miller started with gases such as water vapor, methane, and hydrogen gas, all thought to have been present in the early atmosphere. He put these gases into a sealed flask and then passed a spark through the mixture (Fig. 2.29). On the primitive Earth, lightning might have supplied the energy needed to drive chemical reactions, and the spark was meant to simulate its effects. Analysis of the contents of the flask showed that a number of amino acids were generated.

FIG. 2.29: Could the building blocks of organic molecules have been generated on the early Earth?

BACKGROUND In the 1950s, Earth’s early atmosphere was widely believed to have been rich in water vapor, methane, ammonia, and hydrogen gas, with no free oxygen.

EXPERIMENT Stanley Miller built an apparatus, shown below, designed to simulate Earth’s early atmosphere. Then he passed a spark through the mixture to simulate lightning.

RESULTS As the experiment proceeded, reddish material accumulated on the walls of the flask. Analysis showed that the brown matter included a number of amino acids.

CONCLUSION Amino acids can be generated in conditions that mimic those of the early Earth.

FOLLOW-UP WORK Recent analysis of the original extracts, saved by Miller, shows that the experiment actually produced about 20 different amino acids, not all of them found in organisms.

SOURCE Miller, S. L. 1953. “Production of Amino Acids Under Possible Primitive Earth Conditions.” Science 117:528–529.

Miller and others conducted many variations on his original experiment, all with similar results. Today, many scientists doubt that the early atmosphere had the composition found in Miller’s experimental apparatus, but amino acids and other biologically important molecules can form in a variety of simulated atmospheric compositions. If oxygen gas (O₂) is absent and hydrogen is more common in the mixture than carbon, the addition of energy generates diverse amino acids. The absence of oxygen gas is critical, since Miller-Urey-type reactions cannot run to completion in modern air or seawater. Here, however, geology supports the experiments: Chemical analyses of Earth’s oldest sedimentary rocks indicate that, for its first 2 billion years, Earth’s surface contained little or no oxygen.

Later experiments have shown that other chemical reactions can generate simple sugars, the bases found in nucleotides, and the lipids needed to form primitive membranes. Independent evidence that simple chemistry can form the building blocks of life comes from certain meteorites, which provide samples of the early solar system and contain diverse amino acids, lipids, and other organic compounds.

2.6.2 Experiments show how life’s building blocks can form macromolecules.

From the preceding discussion, we have seen that life’s simple building blocks can be generated under conditions likely to have been present on the early Earth—but can these simple units be stitched together to form the polymeric molecules of life? Once again, careful experiments have shown how polymers could have formed in the conditions of the early Earth. Clay minerals that form from volcanic rocks can bind nucleotides on their surfaces (Fig. 20.30a). The clays provide a surface that places the nucleotides in proximity to one another, making it possible for them to join to form chains, or simple strands of nucleic acid.

Figure 2.30: Spontaneous polymerization of nucleotides.(a) Clays may have played an important role in the origin of life by providing surfaces for nucleotides to form nucleic acids. (b) Addition of a short RNA molecule, the template strand, to a flask containing modified nucleotides results in the formation of a complementary RNA strand.

In a classic experiment, biochemist Leslie Orgel placed a short nucleic acid sequence into a reaction vessel and then added individual chemically modified nucleotides. The nucleotides spontaneously joined into a polymer, forming the sequence complementary to the nucleic acid already present (Fig. 20.30b).

Such experiments show that nucleic acids can be synthesized experimentally from nucleotide building blocks, but until recently the synthesis of nucleotides themselves presented a formidable problem for research on the origins of life. Many tried to generate nucleotides from their sugar, base, and phosphate constituents, but no one succeeded until 2009. That year, John Sutherland and his colleagues showed that nucleotides can be synthesized under conditions thought to be like those on the young Earth. These chemists showed how simple organic molecules likely to have formed in abundance on the early Earth react in the presence of phosphate molecules, yielding the long-sought nucleotides.

Such humble beginnings eventually gave rise to the abundant diversity of life we see around us, described memorably by Charles Darwin in the final paragraph of The Origin of Species:

CHAPTER SUMMARY

• Atoms consist of positively charged protons and electrically neutral neutrons in the nucleus, and negatively charged electrons darting around the nucleus.
• The number of protons determines the identity of an atom.
• Protons and neutrons together determine the mass of an atom.
• Protons and electrons determine the charge of an atom.
• Negatively charged electrons travel around the nucleus in regions called orbitals.
• The periodic table of the elements reflects a regular and repeating pattern in the chemical behavior of elements.

• Electrons that occupy the outermost energy level (shell) of an atom (valence electrons) determine its ability to combine with other atoms to form molecules.
• A covalent bond results from the sharing of electrons between atoms to form molecular orbitals.
• A polar covalent bond results when two atoms do not share electrons equally as a result of a difference in the ability of the atoms to attract electrons, a property called electronegativity.
• A hydrogen bond results when a hydrogen atom covalently bonded to an electronegative atom interacts with an electronegative atom of another molecule.
• An ionic bond results from the attraction of oppositely charged ions.

• Water is a polar molecule because shared electrons are distributed asymmetrically between the oxygen and hydrogen atoms.
• Hydrophilic molecules dissolve readily in water, while hydrophobic molecules in water tend to associate with one another, minimizing their contact with water.
• The pH of an aqueous solution is a measure of the acidity of the solution.
• Water forms hydrogen bonds, which help explain its high cohesion, surface tension, and resistance to rapid temperature change.

• A carbon atom can form up to four covalent bonds with other atoms.
• The geometry of these covalent bonds helps explain the structural and functional diversity of organic molecules.

• Amino acids are linked by covalent bonds to form proteins.
• An amino acid consists of a carbon atom (the α-carbon) attached to a carboxyl group, an amino group, a hydrogen atom, and a side chain.
• The side chain determines the properties of each amino acid.
• Nucleotides assemble to form nucleic acids, which store and transmit genetic information.
• Nucleotides are composed of a 5-carbon sugar, a nitrogen-containing base, and a phosphate group.
• Nucleotides in DNA incorporate the sugar ribose, and nucleotides in RNA incorporate the sugar deoxyribose.
• The bases are pyrimidines (thymine, cytosine, and uracil) and purines (adenine and guanine).
• Sugars are carbohydrates, distinctive molecules composed of C, H, and O atoms, usually in the ratio 1:2:1, that are a source of energy.
• Monosaccharides assemble to form disaccharides or longer polymers called complex carbohydrates.
• Lipids are hydrophobic.
• Triacylglycerols store energy and are made up of glycerol and fatty acids.
• Fatty acids consist of a linear hydrocarbon chain of variable length with a carboxyl group at one end.
• Fatty acids are either saturated (no carbon–carbon double bonds) or unsaturated (one or more carbon–carbon double bonds).
• The tight packing of fatty acids in lipids is the result of van der Waals forces, a type of weak, noncovalent bond.

• In 1953, Stanley Miller and Harold Urey demonstrated that amino acids can be generated in the laboratory in conditions that mimic those of the early Earth.
• Other experiments have shown that sugars, bases, and lipids can be generated in a similar way.
• Once the building blocks were synthesized, they could join together in the presence of clay minerals to form polymers.

Self-Assessment Question 9

Sketch the basic structures of amino acids, nucleotides, monosaccharides, and fatty acids.

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