Building Blocks of Life PDF

Chapter 2: The Building Blocks of Life Contents THE CHEMISTRY OF LIFE........................................................................................................ 2 Elements in biomolecules and biologically relevant ions and trace metals....................................................

Chapter 2: The Building Blocks of Life Contents THE CHEMISTRY OF LIFE........................................................................................................ 2 Elements in biomolecules and biologically relevant ions and trace metals...................................................... 2 Covalent bonds...................................................................................................................................................... 3 Free electron pairs and the shape of molecules.................................................................................................. 5 Hydrogen bonds and the essential properties of water..................................................................................... 7 Van der Waals interactions.................................................................................................................................. 8 Chirality................................................................................................................................................................. 8 Isomerism.............................................................................................................................................................. 9 Functional groups............................................................................................................................................... 10 The redox state of molecules.............................................................................................................................. 11 Important reactivities of functional groups that are common in biology...................................................... 12 Mesomerism........................................................................................................................................................ 12 Acids and bases................................................................................................................................................... 13 CLASSES OF BIOMOLECULES................................................................................................ 14 Lipids, fatty acids, and isoprenoids................................................................................................................... 14 Carbohydrates.................................................................................................................................................... 16 Amino acids and proteins................................................................................................................................... 18 Nucleotides and nucleic acids............................................................................................................................ 19 The Chemistry of Life Chemical principles and transformations are crucial for an understanding of life. The inside of cells contains a densely crowded, complex mixture of biomolecules and inorganic solutes with diverse functions that interact and participate in chemical reactions. For example, enzymes catalyze the multitude of reactions that underly the metabolism of a cell, allowing it to utilize and convert energy from its environment, to produce biomass, and to synthesize the molecules necessary to grow and divide. Chemical catalysts, such as metals, are considered to have played a central role in the transition between abiotic and biotic systems. Still today, metals in enzymes are central to much (but not all) of the biochemistry that takes place within cells and enable the very existence of organisms. The remarkable catalytic properties of enzymes, which we will discuss in Chapter 5, are not achieved by mechanisms special to biology but based on fundamental chemical reactivities. Besides being important for enzymatic reactions, chemical principles also allow cells to form complex, ordered structures, for example for information storage in the case of DNA or to build cytoplasmic membranes that shield the cell interior, among many other functions. Chemical principles were responsible for the emergence of the first organisms and enable powerful applications today. In biotechnology and synthetic biology, an understanding of cellular chemistry is allowing humans to equip organisms with new abilities like generating biofuel from carbon dioxide or digesting plastic waste. So let’s have a closer look at how biology is based on chemistry. Elements in biomolecules and biologically relevant ions and trace metals Currently, 118 different chemical elements are known, of which 94 occur in nature and the remaining are synthetic elements. But biomolecules, and hence living organisms, are mostly composed of only six elements (Table 1): hydrogen (H), carbon (C), nitrogen (N), oxygen (O), phosphorus (P), and sulfur (S). A seventh, less common element found in a few proteins is selenium (Se). Five additional elements, i.e., sodium (Na), potassium (K), magnesium (Mg), calcium (Ca), and chlorine (Cl), are essential and play, for example, important roles in defining the electrochemical membrane potential of cells, and thus their energy balance. Some ions are also relevant as crucial components of active sites in various enzymes, i.e., centers where reactions occur. These include nine so-called trace metals or trace elements: iron (Fe), manganese (Mn), vanadium (V), molybdenum (Mo), wolfram (W), cobalt (Cu), nickel (Ni), copper (Cu), and zinc (Zn). As cofactors of enzymes, they may impart special functions on proteins, such as binding carbon dioxide or transporting molecular oxygen. Table 1: Chemical composition of the bacterium Escherichia coli Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons distributed in shells (or orbitals, depending on the atomic model). Atoms that differ in the number of protons belong to distinct elements. Atoms with the same number of protons, i.e., from the same elements, but different numbers of neutrons are called isotopes. The lightest hydrogen isotope, 1H, is an exception among all atom types because it does not contain neutrons. The periodic table of elements (Figure 1) systematically arranges the diverse elements that exist based on important properties. In the table, the elements are arranged in the order of increasing atomic number (= proton count) from left to right and from top to bottom, starting with hydrogen. The atomic number refers to the number of protons in the nucleus. Vertical columns of the periodic table are termed groups, with elements in a group having the same number of electrons in their outermost shell. This electron number largely influences the chemical properties, such as the number of other atoms an element can connect to. Elements within one group therefore chemically resemble each other. A horizontal row in the table is called a period, and the elements in a period have different properties that change across the row with the number of electrons. Periods correspond to the number of electron shells or orbitals; thus, hydrogen from the first period contains only one shell, while carbon from the second period contains two. Covalent bonds Individual atoms may form chemical bonds to yield molecules. A covalent bond is formed when two electrons are shared between two bonding atoms (Figure 2). The number of covalent bonds that an individual atom may form depends on the position that the element occupies in the periodic table. In neutral (uncharged) molecules, hydrogen (H) forms one bond, carbon (C) four, nitrogen (N) three, and oxygen (O) two. Why is this the case? This information can be obtained from the period table of the elements. As mentioned above, proceeding within a period (i.e., row) stepwise from left to right and starting with one electron in the first group (i.e., column), each element contains one electron more than the previous one. When reaching the right end of the period, the atom contains the maximally possible number of electrons in the electron shell, which is termed the noble gas configuration. In partially filled shells, an electron can form a pair with an electron from another atom to generate a covalent bond. Thus, hydrogen has one bond-forming electron, carbon has four, nitrogen five, oxygen six, etc. When two atoms form a covalent bond, the two paired electrons are shared between them, thus increasing the electron count for each atom. By pairing electrons and thus receiving new electrons from the sharing atom, each atom strives to reach the noble gas configuration with a saturated electron shell. For the first period (upper row) in the periodic table, e.g., for H, saturation is reached with two electrons. Hydrogen, which possesses one electron, can therefore form only one bond by electron-pairing with another atom: 1 own electron + 1 partner electron gives 2 electrons for saturation (Figure 3a). The biologically important elements C, N, and O are located in the second period that can harbor up to eight electrons in the respective shell. Thus, to reach this favorable configuration, they form bonds until the total number of paired electrons around the atom is eight (Figure 3c-d). Carbon atoms therefore form four covalent bonds (4 own + 4 = 8), nitrogen atoms three (5 + 3 = 8), and oxygen atoms only two (6 + 2 = 8). The numbers of bonds an element can form are called valences, i.e., carbon has four valences. Sometimes, atoms share more than one electron with their neighbor to form double or even triple bonds (Figure 4). An example is carbon dioxide (CO2), in which carbon forms a double bond to each of the two oxygen atoms. In this way, atoms can be connected so that each obtains a saturated electron shell, in this case 8 electrons. There are also molecules with triple bonds, which follow analogous electron rules, most prominently molecular nitrogen (N2). Some typical bond lengths in organic molecules are shown in Table 2. Table 2: Common bond lengths in organic molecules Bond type Bond length [pm] C-H 108-110 C-C single bond 146-154 C-C double bond 130-134 C-C triple bond 120 C-N single bond 147-148 C-O single bond 136-147 C-O double bond 116-123 The fact that carbon atoms possess four binding options or valences plays an important role in biology because the high valence number, together with the high stability of C-C bonds (see box “Alternative elements”), provides countless possibilities for binding combinations and hence the diverse architectures of organic molecules. It should be noted that in heavier elements including many metals, the electron configurations are more complex than those described for the first two periods. We will not consider these here but will discuss electrons in metals later in the context of important redox processes.  Alternative elements Elements from the same group in the period table have similar chemical properties. Could life exist that is based on elements other than carbon? Silicon is located below carbon in the table and also has four valences, but it is unlikely that silicon-based life forms have evolved in the universe. Since this element has an additional shell of electrons compared to carbon, silicon atoms are considerably larger. This results in less stable Si-Si bonds and especially double bonds. Silicon- based counterparts of organic molecules might therefore not be suited to support life. In other cases, however, alternative elements are indeed used in biomolecules. Examples are molybdenum and tungsten used as alternative metal cofactors in enzymes, or selenium replacing sulfur in proteins. Representations of molecules You will see molecules represented as structural formulae in several ways, depending on which aspect is important to highlight. Let’s for example consider monosodium glutamate (MSG). This is the sodium salt of glutamic acid, an amino acid that is present as a universal building block in almost all proteins. The ball- and-stick model emphasizes how atoms are connected by covalent bonds (solid lines connecting atoms in Figure 2a). The space-filling model (Figure 2b), on the other hand, shows the differences in atomic size and the spatial arrangement of the atoms in the molecule. The model also includes all the atom connections, but without discriminating between covalent bonds and ionic interactions (more about these in the following sections). Most commonly, molecules are displayed as Lewis structures and their short-hand notation (Figure 2c). These are simplified versions of the ball-and-stick model, and the short-hand notation is abbreviated even further by omitting hydrogen and carbon atoms as the most common atoms in biology. How does one know where which atom is? Carbons are located at all unlabeled corners, branching points, and ends of the structures. The number of hydrogen atoms bound to each carbon is deduced from knowing that carbon forms four bonds in standard molecules. For example, if a carbon has already two bonds to other atoms shown in the structure, the remaining two atoms would be hydrogens. Free electron pairs and the shape of molecules When nitrogen forms three covalent bonds, for example by connecting to hydrogen in ammonia (NH3) molecules, it has a pair of free electrons that does not participate in bonds, visualised as two dots or as a black line near the "N" in Figure 3c. By the same logic, the oxygen atom contains two free electron pairs (Fig. 3d). If one counts the combined number of free electron pairs and bonds around an atom, it is possible to predict the three-dimensional shape of molecules. Since electrons are negatively charged, they can be imagined in a simple model to repel each other and maximise the distance between them. The four covalent bonds in methane thus adopt a tetrahedral geometry. In ammonia and water, corners of the tetrahedron are occupied by free electron pairs and bonds of the central atom. This results in a pyramidal geometry for the ammonia atoms and a bent shape for water (Figure 5). This simple model explains why, e.g., water molecules are not linear although oxygen is bound to two other atoms. For atoms that form double or triple bonds, the same rules apply. CO2 does not contain free electron pairs and is therefore linear unlike the water molecule. We have seen that a set of simple rules based on electron saturation and spatial distancing suffices to not only predict the number of bonds of atoms, but also the geometric shape of the resulting molecules. While these rules explain why water molecules are bent and those of carbon dioxide linear (Figure 4), how does this difference in shape explain that H2O is liquid at room temperature, but CO2 is gaseous although it has a higher molecular mass than H2O? To answer this question, we first need to understand an aspect of molecules that we have not discussed yet: the fact that bond electrons are not evenly distributed between the bonding atoms. The protons in atomic nuclei are positively charged, while electrons have a negative charge, resulting in attractive forces between these two types of particles. Some atomic nuclei exert a stronger pull on the bond electrons than others. This ability of an atom to attract bond electrons is called electronegativity. It is measured on a relative scale with 4 being the highest value f0r fluorine, the most electronegative element that forms bonds (Figure 6a). Electronegativity depends mostly on two factors. First, it increases with the number of positively charged protons in an atomic nucleus. Elements more to the right of the periodic table are therefore more electronegative. Second, electronegativity increases when the bond electrons are closer to the positively charged nucleus, thus elements with fewer shells closer to the top rows of the periodic table are more electronegative. The concept of electronegativity changes our understanding of chemical bonds by providing a continuum of bond types. If two bonding partners have the same electronegativity, they attract the shared electrons equally strongly; the bond is therefore unpolarized (Figure 6b-c). However, two atoms with distinct electronegativity will form polarised bonds in which more electron density is located at the more electronegative element. Since electrons are negatively charged, the more electronegative partner becomes partially negative (denoted as δ-) and the less electronegative partner becomes partially positive (δ+). For very high differences in electronegativity (for example between Na and Cl), both bond electrons are localized at the partner that pulls the strongest (Figure 7). This will result in negative anions and positive cations (Na+ and Cl-). Since no electrons are shared, ions are not linked by covalent bonds, but they maintain strong ionic interactions. These interactions are the characteristic feature of salts, here sodium chloride (NaCl). In the cell, ions, including those of sodium chloride, play important roles as they maintain osmotic balance, generate energy via ion gradients across membranes, and interact with protein structures to stabilize them, among many other functions. Hydrogen bonds and the essential properties of water What is the connection between electronegativity and water being liquid in contrast to CO2? A polarized bond and the resulting partial charges at the atoms can be regarded as an electric dipole. Every dipole has a dipole moment, which can be represented as a vector pointing from the (partially) negative towards the (partially) positive charge site (Figure 8a). Analogous to elementary magnets, the individual dipole vectors of all bonds add up to result in the total dipole moment of a molecule. A water molecule has two angular, polarized bonds that together produce a molecular dipole with the oxygen end being partially negative and the hydrogen end partially positive. Consequently, water molecules are electrostatically attracted to each other and associate via their oppositely polarized regions to form contacts between hydrogen and oxygen. The resulting bond, which is weaker than a covalent bond or an ionic interaction, is called a hydrogen bond (Figure 8b). Hydrogen bonds are the reason for water being liquid at ambient conditions: they hold water molecules together and connect them in extensive, dynamic networks (Figure 8c). The numerous hydrogen bonds impart on water another property that is important for life on Earth. To increase the temperature of liquid water, a large amount of energy is required to first break the numerous hydrogen bonds in the water network until molecules disconnect and can move around freely. Thus, water heats up or cools down more slowly than most other substances, i.e., it has an unusually high heat capacity. Therefore, one important property of water (besides being a solvent as explained further below) is that it acts as a temperature buffer and protects living organisms from the negative consequences of excessive thermal fluctuations. On the other hand, our second example molecule CO2 is linear, has a molecular dipole of zero due to two bond dipole vectors pointing in opposite directions, and is not able to form hydrogen bonds. A small energy input therefore readily separates these molecules to form gaseous CO2, available to carbon-fixing organisms and thus the predominant carbon source of life. As a consequence, a small difference in molecular shape has an enormous impact on the life on this planet. Hydrogen bonds exist in many other molecules. Which general properties do molecules need to have so that these bonds can form? One prerequisite is a hydrogen atom with a sufficiently high positive partial charge, which in biomolecules is the case when hydrogen is bound to oxygen or nitrogen. Second, the other atom participating in the hydrogen bond has to carry a negative partial charge and at least one free electron pair, such as the oxygen in water, but also the oxygen and nitrogen atoms of biomolecules. As we will see later, hydrogen bonds between biomolecules are key to the genetic code and heredity. To a large extent, they are also responsible for the way enzymes and other proteins fold to become functional. Hydrogen bonds strongly influence the solubility of molecules and ions. They are considerably weaker than covalent bonds or ionic interactions (Table 1). Nevertheless, salts like NaCl dissolve in water. This is because dissolved ions are surrounded with a stable sphere of many water molecules (the hydration shell), held together by weak electrostatic forces. The binding energy resulting from numerous electrostatic interactions exceeds the energy of ionic interactions in the salt crystal. Similarly, interactions with water molecules explain why some chemicals are water soluble but others are not. Hydrophilic ("water loving") substances are polarised or ionic, so that water molecules can connect with them via hydrogen bonds and electrostatic interactions to form hydrate shells. Hydrophobic ("water fearing") substances are apolar or weakly polar, hydrogen bonds do not form, and the molecules bind water poorly and aggregate with each other in an aqueous environment. This is an important property of membrane lipids, as will be explained further below. Van der Waals interactions A final example of forces between molecules are the van der Waals interactions (Figure 9a). These are attractive interactions between non-polar molecules. Van der Waals interactions arise when the electron density in a molecule changes to generate a temporary, fluctuating dipole. The partial charges induce opposite dipoles in neighboring molecules, resulting in attractive forces (Table 2.1). These interactions are weak, which may suggest that they are not relevant. However, they are crucial for many important biological phenomena, and many individual molecular contributions can result in large overall effects. Examples in which van der Waals forces play a role are the folding and three-dimensional shape of proteins, the structure of DNA, cellular adhesion, and even the stickiness of gecko feet (Figure 9b). Chirality Let’s consider monosodium glutamate again, which we used to illustrate molecular structures at the beginning of this chapter. We have now described various chemical features and forces that govern molecular properties, with a notable exception: MSG contains a topologically peculiar carbon atom, referred to as a chiral center (alternative terms: asymmetric center or stereocenter). Unlike the other atoms in this molecule, this carbon has four different binding partners (marked with a red circle in Figure 10a). If we recall that the four partners (also termed substituents) point into the corners of a tetrahedron (Figure 10b), the three-dimensional structure can be visualized by a solid wedge for the bond pointing above the paper plane towards you, and a dashed wedge for the bond pointing away from you. A chiral center, i.e., an atom with four different substituents, exists in two mirrored forms, termed configurations, that are not identical to each other. Therefore, the planar chemical formula shown before for MSG represents two different, mirrored molecules, called enantiomers, depending on whether the nitrogen-containing substituent lies in front or behind the carbon chain in the plane of the paper. To move the nitrogen to the opposite side, one would have to rotate the MSG structure horizontally, which would result in two other substituents changing places. Thus, enantiomers are two chiral molecules that are mirrored versions of each other and cannot be interconverted by rotation; they are thus distinct molecules. Conversely, atoms, molecules, or objects that can be mirrored to result in the original topology are called achiral. Besides objects with chiral centers, various other types of chirality exist, including helical chirality as seen in DNA and snail shells (Figure 10c). Another example is human hands, which is why a left glove cannot be put on the right hand. In fact, the word “chiral” is derived from the Ancient Greek word for hand, χείρ (cheir). When two enantiomers, i.e., the mirrored forms, of a chiral object interact with an achiral object, they interact in exactly the same way. For example, one can pick up a ball (achiral) with the left hand just as well as with the right. One enantiomer of MSG has the exact same solubility, melting point, etc. as the other. Why then does chirality matter? The situation is entirely different for interactions between two chiral objects, as everyone quickly realizes who tries to shake hands with the left. Such interactions are ubiquitous in biology because almost all biomolecules including proteins are chiral. Because enzymes discriminate between enantiomers of a molecule in this way, the enantiomers can have vastly different properties in an organism. This became evident, for example, in the thalidomide tragedy. Thalidomide is a chiral substance that was widely used in the 1950s as a sedative and tranquilizer for pregnant women. However, while one of the two thalidomide isomers acts as a tranquilizer (Figure 11), the other can damage a fetus during development. As the prescribed drugs contained both enantiomers (termed a racemic mixture), this had severe consequences. Some molecules contain more than one chiral center. An important example are carbohydrates, such as glucose, which will be described in more detail later in this chapter. Since each chiral center can have two different configurations, the total number of possible molecules with the same overall atom connectivities is 2n, where n is the total number of chiral centers. Pairs of molecule variants are only enantiomers if all their chiral centers are inverted. If molecules differ in only some centers, they are called diastereomers. Unlike enantiomers, diastereomers also exhibit different properties in achiral environments; for example, they can have different melting points or chemical reactivities. Isomerism Enantiomers and diastereomers are examples of substances with the same atom composition but different structures. In general, molecules related in such a way are called isomers (Figure 12). Isomers can also differ in aspects other than chirality. If atoms are connected differently even in two-dimensional representations (for example straight or branched carbon chains), the molecules have a distinct constitution, i.e., they are constitutional isomers. Molecules with the same atom connectivities but different geometrical shapes, such as those with chiral centers, are called stereoisomers. Chirality is not the only phenomenon that results in stereoisomerism. At room temperature, all molecules are subject to rotational movements around single bonds that gives rise to numerous geometrical variants differing in their conformation. These readily interconverting molecule versions are called conformational isomers. Such dynamic processes that change molecular shapes are important in biology, for example in the structural movements of proteins that catalyse enzymatic reactions. Another type of stereoisomerism exists in molecules with carbon-carbon double bonds that, unlike single bonds, do not permit rotations at ambient conditions. Substituents at each carbon can therefore be located on either the same side of the double bond (termed the cis or Z configuration) or on opposite sides (the trans or E configuration). Such stereoisomers that are called cis-trans or E-Z isomers. In daily life in the context of nutrition and health, you might have encountered the term trans-fatty acids or trans-fats as problematic substances in processed foods. These will be explained in Chapter 3 when we discuss lipids. Functional groups Most organic molecules consist of a carbon skeleton, in which individual carbon atoms are linked to each other and to hydrogen atoms (exceptions are single-carbon compounds such as methane). In addition to this core skeleton, they usually also contain moieties termed functional groups (Figure 13). These can be carbon-carbon double bonds (also termed olefinic bonds) or triple bonds, or, more commonly, atoms other than C and H, the heteroatoms (in biology most commonly O, N, P, and S). Many biomolecules feature more than one functional group. Functional groups often define the molecular properties of the entire compound. The presence of heteroatoms in functional groups is often associated with uneven electronic distributions and therefore partial charges and dipole moments in the molecule. These intramolecular imbalances determine how a molecule interacts with its chemical environment, with consequences for physical properties such as melting and boiling points influenced by molecular adhesion, or for the way how one molecule reacts with another. Furthermore, the presence of functional groups that can form hydrogen bonds or engage in other interactions with water influences whether a molecule is water-soluble or not. Particularly common in cellular metabolites are oxygen-containing functional groups, such as alcohol, aldehyde, keto, and carboxyl groups. The corresponding substance classes are called alcohols, aldehydes, ketones, and carboxylic acids, respectively (Figure 14). While representatives of the first three classes are usually uncharged, carboxylic acids tend to form anions under physiological conditions that strongly interact with water, (metal) cations, or other polar molecules. For more information on acidity as an important feature, see the section “Acids and Bases” below. Nitrogen is predominantly found in amino and amide groups or incorporated into ring structures (Figure 15). The term primary, secondary, and tertiary amine (and, analogously, also amide) refers to the number of carbon-containing substituents that are attached to the nitrogen atom in the functional group. Proteins, for example contain numerous secondary amide moieties, also referred to as peptide bonds, each of which connects two amino acids, (see Chapter 3). Phosphorous is almost exclusively incorporated as mono-, di-, or triphosphate groups into biological molecules (Figure 16). One of the most important molecules in biology is adenosine triphosphate (ATP), in which the three phosphate groups form two anhydride bonds (between phosphate units) and one ester bond (between phosphate and carbon). Cleavage of such bonds to free phosphate, which is a highly stable compound, is associated with a large decrease in free energy. Attachment of phosphate or the adenosine monophosphate (AMP) portion of ATP to biomolecules therefore activates these compounds for follow-up reactions and thus enables diverse biochemical processes that would otherwise be endergonic, e.g., in the synthesis of macromolecules such as proteins. For more details on the central importance of ATP see Chapter 6. Functional groups containing sulfur atoms (Figure 16) are named in analogy to functional groups containing oxygen: Thiols correspond to alcohols, and thioethers and thioesters are the sulfur analogs of ethers and esters, respectively. Disulfide groups contain two sulfur atoms linked to each other. Disulfide bridges are common in proteins, where they play an important role in stabilizing the three-dimensional shape of these large biomolecules via covalent crosslinks. Thioesters such as acetyl- CoA contain high-energy carbon-sulfur bonds that are easily cleaved, a principle that cells often use to activate a carboxylic acid, in this case acetic acid, for subsequent reactions. The redox state of molecules Biomolecules contain a wide variety of functional groups, which can be classified depending on the heteroatoms present in the group. Since many heteroatoms are more electronegative than carbon atoms, they attract the electrons of the covalent bond. For example, the carbon in a carboxylic acid group has more carbon-oxygen bonds than in aldehyde or alcohol groups. CO2 is therefore more highly oxidized than formic acid, formaldehyde, methanol, and methane (Figure 17a). Vice versa, the carbon in methane is the most highly reduced in this series of C1 compounds, while that in CO2 is the least reduced. The degree to which an atom is oxidized (or reduced) is termed its redox state, which can be described in terms of oxidation numbers assigned to atoms in a molecule (Figure 17b). Oxidation numbers are identified by formally assigning both electrons of a polarized bond to the more electronegative partner and determining the hypothetical resulting charge of the atoms. In a carbon- oxygen bond, for example, both electrons are assigned to oxygen, providing a formal charge of +1 to the electron-deficient carbon and of -1 to oxygen. CO2 has four carbon- oxygen bonds (from two double bonds), resulting in a total oxidation number of +4 for carbon and -2 for each oxygen. In CH4, carbon is less electronegative than hydrogen, giving an oxidation number of -4 for carbon and +1 for each hydrogen. Following these rules, you can calculate the oxidation numbers for the carbon in formic acid (+2), formaldehyde (0), and methanol (-2). While the word oxidation might suggest that oxygen is always involved, this is not the case, as the redox rules can be applied to any molecule and atom, including metal ions. For example, the nitrogen equivalent of alcohols regarding the carbon oxidation number are amines with a C-N bond, a nitrogen-containing counterpart of aldehydes are imines (or Schiff bases) with a C=N double bond, and an equivalent of carboxylic acids are amides (Figure 13). In general, oxidation refers to a chemical process that increases the oxidation state by removal of electrons, and reduction refers to a gain in electrons. As we have seen in Chapter 1, such changes in redox states are important because life is fueled by redox processes, in which electrons are passed from reduced donor to oxidized acceptor molecules. In methanogenesis, which probably was one of the oldest modus operandi of early life forms, these electrons are stripped from molecular hydrogen and transferred to carbon dioxide, thereby transforming the most oxidized one-carbon molecule (after several further steps) into methane, the most reduced C1 compound. We will see in Chapter 6 how this transformation is used to sustain the metabolic activities of living organisms. Important reactivities of functional groups that are common in biology Depending on the distribution of partial charges and the presence or absence of free electron pairs, functional groups show characteristic behavior in chemical reactions (Figure 18). Many of these reactivities can be understood or even predicted by taking into account that electron-deficient parts of molecules preferably react with electron- rich ones, the latter often in the form of free electron pairs. A reaction type rationalized by this concept is a redox reaction, in which electrons are transferred from the electron-rich to the electron-poor molecule, with the latter being reduced and the former being oxidized in the process. In other reactions, the redox state of molecules does not change. A reaction of major importance in biology is nucleophilic substitution, in which a group with free electrons, termed the nucleophile (“the nucleus lover”), attacks a partially positive atom and displaces another functional group, the leaving group, that is bound to this atom. This is, for example, the principle in phosphate transfer reactions (Chapter ###), in which a phosphate group or larger phosphate-containing portion of ATP is attached to molecules to activate them for subsequent reactions. If a nucleophilic substitution occurs at the carbonyl carbon of a carboxylic acid, amide, ester, thioester, or related moiety, it is called nucleophilic acyl substitution. The term “acyl” refers to the moiety comprising the carbonyl group and its carbon substituent. An important example is the formation of peptide (i.e., primary amide) bonds in protein biosynthesis (Chapter ###), another is the substitution at thioester carbons in fatty acid biosynthesis (Chapter ###). In the hydrolysis of proteins, water acts as the nucleophile to disconnect peptide bonds, resulting in protein cleavage. In all of these cases, the carbon has a leaving group with a hetero atom (O, N, S) that can be displaced by the nucleophile. For aldehydes and ketones that do not contain such leaving groups, nucleophiles attack but do not displace another group. Here, the double bond is broken, and the nucleophile is added to the carbon. This reaction is called nucleophilic addition. We will encounter it below in the context of carbohydrates such as glucose, which contains the nucleophile (in this case an OH group) and an aldehyde moiety in the same molecule, resulting in ring formation. If addition occurs on a non-polarized olefinic bond instead of a carbonyl group, it is facilitated when the initially attacking reactant is electron- deficient and interacts with the electron-rich double bond system. Such a reactant is an electrophile, i.e., an “electron-lover”, and the reaction is called electrophilic addition. These are just some of the many reaction types that exist in cellular biosynthetic pathways, and we will encounter further examples when these pathways are discussed in more detail. Mesomerism So far, we imagined bonds as entities that remain fixed in molecules. However, in some cases the additional electron pair in double bonds, termed π-electrons, can distribute over larger molecule regions or even entire molecules. This happens when double bonds alternate with single bonds along a series of atoms, a situation termed conjugation (Figure 19). Conjugated double bonds are not localized but dispersed as an electron cloud over the entire conjugated system, they are delocalized electrons. This phenomenon is known as mesomerism. Electron delocalization is not directly obvious from the Lewis structures (see Box ###), which are used typically to depict molecules. One way to visualize mesomerism is to show various versions of the same molecule, in which electrons are flipped from one position to the next (Figure 19). These structures are called resonance structures. It is important to realize that electrons do not actually change positions and that the resonance structures do not show real molecules or intermediates but rather help explain how molecules might react. Another representation that is closer to reality is to draw delocalized electron pairs as dotted lines that extend over the mesomeric portion of the structure. Why is mesomerism important? For example, mesomerism can also delocalize charges if these are positioned in conjugation to a π-electron system, resulting in more stable molecules, in which single atoms do not carry a full charge anymore. Likewise, free electron pairs can participate in mesomerism. This increased stability through mesomerism influences many important properties, one of them being acidity explained in the next section. Acids and bases An acid is a substance that can donate a hydrogen cation, also called a proton (H+). As a result, the remainder of the molecule becomes anionic (Figure 20). In general, a separation of molecular components into ions is termed dissociation. The acidity of a molecule correlates with the ease with which this process occurs. In biomolecules, dissociation and reassociation of protons and anions occur constantly. Because of this equilibrium, only a fraction of the total molecules is therefore dissociated. The stronger the acid, the larger this fraction of dissociated molecules is. If a molecule (for example the acid anion) can bind a proton, it is called a base. Often, representations of biochemical reactions involving acids show protons, but it is important to realize that free protons are not stable in aqueous environments because the process of proton donation is always coupled with proton uptake (protonation) by another molecule in an acid-base reaction. The most abundant compound in the cell that acts as a proton acceptor, i.e., as a base, is water. The protonation of water, in which one free electron pair of oxygen forms a bond with the proton, results in a hydronium cation (H3O+). However, even if hydronium ions are the main protonation product, reaction schemes often show free protons for simplification. Water can not only act as a base, but also as an acid. The molecular species resulting from proton donation and proton uptake in water are a hydroxide anion (OH-) and a hydronium cation (H3O+), respectively. Water is only weakly acidic: only one in ca. 550 million water molecules is dissociated under physiological conditions. However, much stronger acids in biological systems are carboxylic acids, which readily dissociate to form carboxylate anions and hydronium cations if water is the proton acceptor. Facile deprotonation can be explained by mesomerism (see the previous section): The resulting carboxylate moiety features a carbonyl π-electron pair in conjugation with free electron pairs at the anionic oxygen (Figure 20). Via mesomerism, the negative charge can be distributed among two oxygen atoms, thus stabilizing the ion and rendering it more stable. Similarly, mesomerism stabilizes free phosphate groups, an effect that is important for reactions involving phosphorylation and dephosphorylation in the context of ATP (see Chapter ###. Various ways exist to quantify acidity. The acidity of an aqueous solution can be expressed using pH values, which is defined as the negative logarithm of the H3O+ concentration. These concentration values cover a huge range, and the negative logarithm is therefore used to convert them to a convenient scale from around 0 to 14, with values below 7 for acidic solutions and above 7 for basic ones. pH values only provide information on a solution, not on the property of the dissolved acid or base itself. For example, a weak acid at high concentration can result in the same pH as a strong acid at low concentration. To provide information on the acidity of a molecule, pKa values are used, which corresponds to the tendency to donate a proton, more specifically to the pH at which half of the molecules are deprotonated. Again, lower values indicate stronger acids. Classes of biomolecules Having looked at the chemical elements and functional groups present in biomolecules, let’s consider the molecules themselves. In the central metabolism of all organisms, four large classes of biological molecules are prevalent: lipids, carbohydrates, amino acids, and nucleotides. The latter three groups, but not the lipids, also provide building blocks for macromolecules that are formed by covalent concatenation of single units. Some macromolecules even combine building blocks of different monomer types. These macromolecules differ greatly in their chemical composition, and they also serve distinct purposes that include the formation of cellular structures, energy storage, information storage and transfer, and biocatalysis. However, they share a common biosynthetic principle termed condensation (Figure 21). The term condensation refers to the loss of water during a reaction. However, the removal of water molecules only applies to the net chemical equation; in reality, the reactions are more complicated and involve a sequence of biochemical steps involving activation of a reaction partner. Often, the oxygen is activated by linkage to a phosphate group, AMP, or a similar unit, enabling its connection to another molecule. In the following sections, we will take a closer look at the biomolecule classes. Lipids, fatty acids, and isoprenoids Lipids are a heterogeneous class of compounds that are grouped together not because of a common biochemical origin but because of their unifying property of being soluble in fat, but poorly soluble in water (Figure 22). Therefore, oil, fat, and wax can also be classified as lipids. In bacteria and eukaryotes, fatty acid-based lipids are predominant, but archaea, in contrast, contain lipids based on distinct types of compounds called isoprenoids or terpenes. Fatty acids are composed of long hydrocarbon chains that have a carboxylic acid group at one end. The carboxyl carbon is denoted as the first carbon atom of the chain. Usually, fatty acids are incorporated into larger molecules that are used by organisms for energy storage, formation of cell membranes, and other important functions. Depending on whether the hydrocarbon chain contains double bonds or not, fatty acids are subclassified as saturated (lacking double bonds) or unsaturated (containing one or more double bonds) (Figure 23). These double bonds are predominantly Z- configured, which results in bent chains. Such differences in molecular shape are important: Margarine, for example, primarily contains saturated fatty acid chains that align well and are tightly packed together. Olive oil, on the other hand, contains a high proportion of unsaturated fatty acid moieties that prevent such tight packing. Therefore, olive oil is liquid at room temperature, while margarine is solid. Related phenomena are used by organisms to control the fluidity of cell membranes (for more details, see Chapter ###). Lipids containing E-configured fatty acids are not common in nature but can result from double bond isomerization in synthetic food processing. These compounds represent a health risk and are monitored in many countries. The length of the fatty acyl hydrocarbon chain is inversely correlated with water solubility. Acetic acid, a C2 compound, is infinitely miscible with water because the polar carboxylic acid group constitutes a major portion of the molecule. The same is true for the C1 acid formate. With increasing chain length, the hydrophobic character of the hydrocarbon moiety gains importance and molecules become poorly soluble in water and tend to aggregate via their lipophilic portions. This has important consequences for biology. Long-chain fatty acids such as the C16 compound palmitic acid, the most abundant fatty acid in bacterial cells, contain two molecular regions with opposite properties. The carboxylic acid group is hydrophilic and easily forms hydrogen bonds with water or other carboxylic acid groups, while the hydrocarbon chain is hydrophobic (or lipophilic). The entire molecule combining both features is called an amphiphile, from Ancient Greek ἀμφί (amphi), meaning “on both sides”. In water, fatty acids arrange themselves into sphere-like structures called micelles (Figure 24). These microdroplets are filled with the hydrophobic hydrocarbon chains of the component fatty acids. Their hydrophilic carboxylic acid groups are located on the micelle surface and form hydrogen bonds with the surrounding water. Micelles form spontaneously when washing with soap: the amphiphilic soap molecules intermingle with the lipids in grease and dissolve it by forming small droplets, which can be rinsed away. Micelles are excellent hosts for hydrophobic substances, but water-based reactions cannot take place in their lipid interiors. Much more important for biology is the fact that fatty acid-containing amphiphilic molecules, such as phospholipids, can also form higher-order, water-filled compartments called vesicles (Figure 24). Their amphiphilic nature enables them to form a lipid double layer or membrane, which constitutes a hydrophobic barrier separating an aqueous chamber from its environment. This compartmentation is a key condition for life. Biological membranes are not composed of free fatty acids, but mostly of phospholipids, which render the membrane bilayers more stable: In contrast to fatty acids, in which the carboxyl group is only partially dissociated into ions at physiological conditions, phospholipids have permanently charged head groups that serve as a powerful hydrophilic moiety. This group is linked to two fatty acids, increasing the hydrophobicity and changing the overall molecule shape to a conical geometry that influences the aggregation properties. Cellular membranes usually contain a mixture of saturated and unsaturated fatty acids, of which the latter render the membrane more fluid as explained above. In addition to fatty acids, isoprenoids (also called terpenoids) are a second, distinct class of lipid building blocks. These feature carbon chains that usually contain methyl branches at every fifth position, in contrast to fatty acids that are commonly unbranched. A major difference between bacteria and eukaryotes on the one hand and archaea on the other is that archaeal membranes are uniquely based not on fatty acids but on isoprenoids. It is therefore difficult to tell how the membrane of LUCA, the last universal common ancestor, looked like. However, isoprenoids are common in all organisms. An important isoprenoid lipid in many animals including humans is cholesterol, which is embedded in the fatty acid-based cell membrane to increase fluidity. In addition, a vast diversity of other substances exist that are isoprenoid- based and serve a multitude of functions, including redox coenzymes, cellular protectants, lipophilic membrane anchors for proteins, hormones, and light sensors. Carbohydrates Carbohydrates are often referred to as saccharides or sugars, but not all members of this class of molecules taste sweet; some, including cellulose, have no taste at all. The name carbohydrates is based on the fact that many of these compounds have atomic compositions that can be formally written as Cn(H2O)n. This is, however, misleading, as carbohydrates are not formed by reaction of carbon with water or contain water molecules, but are the product of complex biosynthetic pathways. Carbohydrates can have vastly diverse structures ranging from small compounds with less than 10 carbon atoms termed monosaccharides (Figure 25) to macromolecules (polysaccharides) with more than 2000 carbons, in which numerous monosaccharides are linked together (Figure 26). Saccharides of intermediate complexity are called disaccharides if they contain two monomers and oligosaccharides for typically 3-10 monomers. Likewise, carbohydrates can be part of other biomolecules including lipids and proteins. We will first take a closer look at the monosaccharide building blocks before discussing the more complex biomolecules. Monosaccharides consist of carbon chains of different lengths that carry hydroxyl groups on most of the carbon atoms and, in addition, either an aldehyde group at one end, denominated as C1, or a keto group at C2 (Figure 25). Aldehyde carbohydrates, e.g., glucose, are termed aldoses and those with a keto group, such as fructose, ketoses. The most widespread monosaccharide is glucose with 6 carbons, an aldose with hydroxyl groups on C2-C6. Glucose or its phosphorylated version glucose-6- phosphate is the most important source of energy in many living organisms. Glucose is also known as dextrose and tastes sweet. Note, however, that the household sugar that is more common in daily life is a disaccharide termed sucrose that structurally differs from glucose. Sucrose is widespread in plants, among other organisms, and is usually extracted from sugar cane or sugar beet. In addition to C6-carbohydrates, termed hexoses, molecules with shorter or longer chain lengths but similar structural features exist. The smallest aldose is a C3 compound (i.e., a triose) named glyceraldehyde. The 3-phosphorylated form glyceraldehyde-3-phosphate occupies a central place in central carbon metabolism, as we will see in Chapter 6. Important carbohydrates with five carbon atoms, termed pentoses, include ribose, a building block of ribonucleic acid (RNA). In contrast, the pentose in desoxyribonucleic acid (DNA) is 2-desoxyribose, a derivative of ribose that carries a hydrogen instead of a hydroxyl group at the C2 atom. Because it lacks an oxygen atom, its molecular formula deviates from the generic Cn(H2O)n. As for glucose and glyceraldehyde, ribose and desoxyribose are commonly present in cells as forms in which the terminal C atom, in this case C5, is phosphorylated. Carbohydrates feature several structural peculiarities that contribute to their high diversity. In water, monosaccharides with five or more carbon atoms exist in several different chemical forms that continuously interconvert: one open-chain form with a carbonyl group and several closed-chain hemiacetal forms (Figure 25b). These interconversions arise from thermal molecular motion that makes the carbohydrate chain swing around. The hemiacetal structures result from bond formation between a nucleophilic hydroxyl group and the electrophilic carbonyl carbon, i.e., the aldehyde group of an aldose or the keto group of a ketose, resulting in hemiacetal formation. Depending on the way the ring closes, the hemiacetal hydroxyl group in the cyclic form can point either down or up. The former molecule is called the α form, the latter the β form. However, hemiacetals are easily opened again when surrounding water molecules interact with the molecule by forming new hydrogen bonds. Although all hydroxyl groups of carbohydrates are nucleophilic, small rings below 5 and above 6 atoms are unstable and are therefore not formed. Glucose forms almost exclusively 6- membered rings, while the preferred ring size of pentoses such as ribose is 5. A second source of structural diversity in carbohydrates is based on chirality. As discussed earlier in this chapter, carbon atoms carrying four different substituents are chiral and can exist in two different, non-interconvertible forms (configurations). In carbohydrates, this is the case for all carbons with hydroxyl groups except for the terminal carbon, for example C2-C5 in glucose. The presence of multiple chiral carbons, each of which can have two configurations, results in high chiral complexity with many different stereochemical variants. Glyceraldehyde, a triose with one chiral center, forms two stereoisomers, while 24 = 16 stereoisomers exist for C6 aldoses such as glucose with four chiral centers. To facilitate the comparison between such forms, monosaccharides are often shown in a representation called the Fischer projection (Figure 27). By convention, the structure is drawn as a straight carbon chain with the highest-oxidized carbon at the top, for example the aldehyde group in an aldose. Horizontal bonds are imagined as pointing towards the observer. What is the naming system for the numerous carbohydrates? All compound pairs that feature the same relative configuration, i.e., a carbohydrate and its mirrored, enantiomeric form, have a distinct chemical name. For example, the hexoses glucose and galactose (Figure 27) are named differently because they feature a distinct relative orientation of hydroxyl groups. On the other hand, “glucose” is used for the two enantiomeric compounds with the same relative configuration. To distinguish a carbohydrate from its enantiomer, the denominators L (from Latin laevis: left) and D (from dexter: right) are added to the name, which refer to the orientation of the hydroxyl group at the second-lowest carbon (the lowest chiral center) in the Fischer projection. In cells, almost all biological carbohydrates, including glyceraldehyde, ribose, glucose, and galactose, are present in the D-configuration. This is remarkable, since abiotic processes that preceded life should have resulted in equal amounts of mirrored stereoisomers. What happened during early evolution that created this break of chiral symmetry? Various hypotheses are proposed to explain this phenomenon, but the answer remains unknown. An important consequence is that chiral compounds do not interact equally well with both enantiomers of another biomolecule. For example, organisms cannot utilize L-glucose as a source of energy. A small enantiomeric excess in a compound present in a catalytic system might therefore have rapidly induced a chiral symmetry break in other substances. As mentioned above, monosaccharides are often linked together to result in oligo- or polysaccharides (Figure 26). The products contain the building blocks as cyclic forms connected via two hydroxyl groups by formal loss of water. In reality, the attachment of a carbohydrate unit, termed glycosylation, is more complex and requires the chemical activation of one of the hydroxyl groups. This is usually the hemiacetal hydroxyl group of the cyclized carbohydrate. While hemiacetals are unstable in water and reversibly form the open-chain carbohydrate, the acetal resulting from glycosylation is stable. Therefore α and β configurations also remain fixed. By glycosylation, carbohydrate mono- and oligomers can also be attached to a wide range of other biomolecules including proteins, lipids, and other metabolites. This process serves a wide range of important purposes including protein folding, cell-cell adhesion, regulation, immune recognition, pathogen evasion, and further phenomena. Amino acids and proteins Amino acids are metabolites with two functional groups: an amino and a carboxylic acid group. While a broad diversity of amino acids exist, 20 of them are the building blocks of proteins. A unifying feature of these so-called proteinogenic amino acids is that the two functional groups are bound to the same carbon atom. This carbon is called α because it is the first carbon adjacent to the carboxyl group, and proteinogenic amino acids are therefore α-amino acids (Figure 28). With the exception of glycine, the simplest amino acid, all other amino acids carry a further substituent and thus feature a chiral α-carbon. All of these proteinogenic amino acids belong to the L-stereoisomers, in which the amino group is situated on the left (laevis in Latin) in the Fischer projection (Figure 29). The lack of D-amino acids in proteins is another case of broken chiral symmetry, a situation that we have already encountered for carbohydrates above. In organisms, amino acids can be connected to form products termed peptides for chains of 2-100 building blocks (also termed residues), or proteins if the chains are larger. To generate such compounds, the carboxylic acid group of one amino acid would be formally connected to the amino group of the next unit under loss of water to form a primary amide bond, the peptide bond (Figure 30). However, as for carbohydrates, the actual biochemical process does not involve free amino acids but is more complex and will be explained in Chapter ###. For conventional peptide and protein biosynthesis, cells use ribosomes, highly efficient nanomachines that successively incorporate amino acid residues into the growing molecule chain, all connected by peptide bonds. The different side chains endow the 20 proteinogenic amino acids with distinct chemical and structural properties. Based on these side chains, amino acids can be roughly grouped into non-polar, polar, negatively charged (or basic), and positively charged (or acidic) compounds (Figure 28). The side chains confer various functional characteristics to proteins and define their three-dimensional structure and polarity, as well as the chemical reactivity in case of enzymes. But even though the functional repertoire of amino acid side chains is broad, the side chains cannot accomplish all reaction types. Electrophilic side chains, for example, do not exist among the 20 amino acids. For some reactions, proteins therefore depend on cofactors as functional enablers. Some cofactors are organic molecules and then called coenzymes. One example among the many known coenzymes is heme and related compounds that bind a metal atom in their center. Heme contains iron and facilitates redox processes or bind oxygen; another heme variant contains magnesium at its center and participates in light-harvesting electron shuttling through photosynthetic reaction centers in phototrophic bacteria and chloroplasts. Other cofactors are inorganic and metal-based. Some of these consist of iron-sulfur clusters that are reminiscent of the minerals found, for example, in the walls of the bubbles in white smokers (see Chapter 1). These structures stand out for their useful properties in redox reactions: They easily accept and donate electrons. Nature goes to great lengths to synthesize such mineral-like clusters which still today play central roles in many redox reactions. It is tempting to suspect the iron-sulfur clusters in enzymes as remnants of ancient abiotic processes that are still preserved in biological systems. Nucleotides and nucleic acids Nucleotides are the basic building blocks of nucleic acids, such as DNA and RNA, which constitute the hereditary and information-encoding material in all organisms. In addition, nucleotides also are integral parts of various coenzymes and other molecules involved in enzymatic reactions. Free nucleotides also provide chemical energy in cells to propel energy-demanding processes. An important example of a nucleotide is ATP (Figure 31a). Its turnover in living organisms is impressive: humans produce and use an amount of ATP every day that approximately sums up to their own body weight. All these fundamental functions of nucleotides suggest that they belonged to the earliest molecules of life. Nucleotides have a modular structure and are composed of three different building blocks: a nucleobase, a pentose that is either D-ribose or D-2-deoxyribose, and one or more phosphate units (Figure 31b). Nucleotides can feature between one and three phosphate groups linked together via anhydride bonds and are accordingly called mono-, di-, or triphosphates. Some nucleotide variants do not contain phosphate and are then called nucleosides. In DNA and RNA, which have similar but not identical structures, nucleotides are linked to long chains via single phosphates that each connect the pentose units of two building blocks. The pentose in DNA is D-2- deoxyribose, while RNA contains D-ribose. In both DNA and RNA, the monophosphate bridge is positioned between carbon 3 of one pentose and carbon 5 of the other. When speaking of pentoses in nucleotides and nucleic acids, these positions are named 3’ (read as “three-prime”) and 5’ (“five-prime”). The reason is that the nucleobases attached to the ribose have their own ring numbering system starting with 1. To avoid confusion, pentose numbers are therefore marked with the prime symbol. The 3’ and 5’ positions are particularly important because they also occur at the opposite ends of nucleic acid molecules and are, for example, used when describing the direction of a DNA or RNA chain or its growth during polymerization. How can information be encoded in these macromolecules? This “code” is realized by the succession (i.e., the sequence) of different nucleobases attached to the sugar backbone. The sequence is composed of four different nucleobases present in either DNA or RNA (Figure 32). Comparing DNA to RNA, three of these nucleobases are present in both nucleic acids: adenine, guanine, and cytosine. The fourth nucleobase is unique and consists of thymine in DNA and uracil in RNA. Nucleobases are either monocyclic and then termed pyrimidine bases (cytosine, thymine, and uracil) or bicyclic and named purine bases (adenine and guanine). They all contain nitrogen atoms in their rings, one of which is the attachment point to the pentose. A characteristic feature of DNA is that two individual chains, or strands, pair via hydrogen bonds between nucleobases and adopt a helix-type structure, the DNA double helix. In contrast to the double-stranded DNA, RNA is single-stranded. Base- pairing via hydrogen bonds is highly specific: adenine always pairs with thymine, whereas guanine always pairs with cytosine. Through their complementary sequences, versions of the same code are stored in each DNA strand, which is, for example, important in distributing the genetic information to the two daughter cells during cell division. More information on these and other processes in DNA and RNA that are important for the evolution and existence of life will be given in Chapter 4.