《生物化学》课程教学资源（书籍文献）Biochemistry, EIGHTH EDITION 2015, Jeremy M. Berg, John L. Tymoczko, Gregory J. Gatto, Jr. Lubert Stryer.pdf_P36-P40

6 CHAPTER 1 Biochemistry: An Evolving Science 1.3 Concepts from Chemistry Explain the Properties of Biological Molecules We have seen how a chemical insight into the hydrogen-bonding capabilities of the bases of DNA led to a deep understanding of a fundamental biological process. To lay the groundwork for the rest of the book, we begin our study of biochemistry by examining selected concepts from chemistry and showing how these concepts apply to biological systems. The concepts include the types of chemical bonds; the structure of water, the solvent in which most biochemical processes take place; the First and Second Laws of Thermodynamics; and the principles of acid–base chemistry. The formation of the DNA double helix as a key example We will use these concepts to examine an archetypical biochemical process— namely, the formation of a DNA double helix from its two component strands. The process is but one of many examples that could have been chosen to illustrate these topics. Keep in mind that, although the specific discussion is about DNA and double-helix formation, the concepts considered are quite general and will apply to many other classes of molecules and processes that will be discussed in the remainder of the book. In the course of these discussions, we will touch on the properties of water and the concepts of pK a and buffers that are of great importance to many aspects of biochemistry. The double helix can form from its component strands The discovery that DNA from natural sources exists in a double-helical form with Watson–Crick base pairs suggested, but did not prove, that such double helices would form spontaneously outside biological systems. Suppose that two short strands of DNA were chemically synthesized to have complementary sequences so that they could, in principle, form a double helix with Watson–Crick base pairs. Two such sequences are CGATTAAT and ATTAATCG. The structures of these molecules in solution can be examined by a variety of techniques. In isolation, each sequence exists almost exclusively as a single-stranded molecule. However, when the two sequences are mixed, a double helix with Watson–Crick base pairs does form (Figure 1.8). This reaction proceeds nearly to completion. What forces cause the two strands of DNA to bind to each other? To analyze this binding reaction, we must consider several factors: the types of interactions and bonds in biochemical systems and the energetic favorability of the reaction. We must also consider the influence of the solution conditions—in particular, the consequences of acid– base reactions. Covalent and noncovalentbonds are important for the structure and stability of biological molecules Atoms interact with one another through chemical bonds. These bonds include the covalent bonds that define the structure of molecules as well as a variety of noncovalent bonds that are of great importance to biochemistry. Covalent bonds. The strongest bonds are covalent bonds, such as the bonds that hold the atoms together within the individual bases shown on page 4. A covalent bond is formed by the sharing of a pair of electrons between adjacent atoms. A typical carbon–carbon (C}C) covalent bond has FIGURE 1.8 Formation of a double helix. When two DNA strands with appropriate, complementary sequences are mixed, they spontaneously assemble to form a double helix. G T T T A A A C G C A A A T T T G T T T A A A G C A A A T T T C

a bond length of 1.54 Å and bond energy of 355 kJ mol 1 (85 kcal mol 1 ). Because covalent bonds are so strong, considerable energy must be expended to break them. More than one electron pair can be shared between two atoms to form a multiple covalent bond. For example, three of the bases in Figure 1.6 include carbon–oxygen (C“O) double bonds. These bonds are even stronger than C}C single bonds, with energies near 730 kJ mol 1 (175 kcal mol 1 ) and are somewhat shorter. For some molecules, more than one pattern of covalent bonding can be written. For example, adenine can be written in two nearly equivalent ways called resonance structures. N N N N H H NH2 5 4 N N N N H H NH2 5 4 These adenine structures depict alternative arrangements of single and double bonds that are possible within the same structural framework. Resonance structures are shown connected by a double-headed arrow. Adenine’s true structure is a composite of its two resonance structures. The composite structure is manifested in the bond lengths such as that for the bond joining carbon atoms C-4 and C-5. The observed bond length of 1.40 Å is between that expected for a C}C single bond (1.54 Å) and a C“C double bond (1.34 Å). A molecule that can be written as several resonance structures of approximately equal energies has greater stability than does a molecule without multiple resonance structures. Noncovalent bonds. Noncovalent bonds are weaker than covalent bonds but are crucial for biochemical processes such as the formation of a double helix. Four fundamental noncovalent bond types are ionic interactions, hydrogen bonds, van der Waals interactions, and hydrophobic interactions . They differ in geometry, strength, and specificity. Furthermore, these bonds are affected in vastly different ways by the presence of water. Let us consider the characteristics of each type: 1. Ionic Interactions . A charged group on one molecule can attract an oppositely charged group on the same or another molecule. The energy of an ionic interaction (sometimes called an electrostatic interaction) is given by the Coulomb energy: E 5 kq1q2/Dr where E is the energy, q1 and q2 are the charges on the two atoms (in units of the electronic charge), r is the distance between the two atoms (in angstroms), D is the dielectric constant (which decreases the strength of the Coulomb depending on the intervening solvent or medium), and k is a proportionality constant (k 5 1389, for energies in units of kilojoules per mole, or 332 for energies in kilocalories per mole). By convention, an attractive interaction has a negative energy. The ionic interaction between two ions bearing single opposite charges separated by 3 Å in water (which has a dielectric constant of 80) has an energy of 2 5.8 kJ mol 1 ( 2 1.4 kcal mol 1 ). Note how important the dielectric constant of the medium is. For the same ions separated by 3 Å in a nonpolar solvent such as hexane (which has a dielectric constant of 2), the energy of this interaction is 2 232 kJ mol 1 ( 2 55 kcal mol 1 ). q1 q2 r Distance and energy units Interatomic distances and bond lengths are usually measured in angstrom (Å) units: 1 Å 5 10210 m 5 1028 cm 5 0.1 nm Several energy units are in common use. One joule (J) is the amount of energy required to move 1 meter against a force of 1 newton. A kilojoule (kJ) is 1000 joules. One calorie is the amount of energy required to raise the temperature of 1 gram of water 1 degree Celsius. A kilocalorie (kcal) is 1000 calories. One joule is equal to 0.239 cal. 7 1.3 Chemical Concepts

8 CHAPTER 1 Biochemistry: ges on nearb y atoms attracting another Hy An Evolving Scienc hydrogen atom in a hydrogen bond is partially shared by two electronega- tive atoms such as nitrogen or oxygen.The hydrogen-bond donor is the group ond omorpto that inclu both the atom to whic ed and the ch the hydrogen atom is more tightly whereas the 一#… 1.9 ptor i N- -0 ich the onde t he fro tia 0-H-------N positive charge().Thus,the hydrogen atom with a partial positive charge 0-H 0 can interact with an atom having a partial negative charge()through an ionic interaction. and are shown. Ri Hydrogen bonds are much weaker than covalent bonds.They have ener k mo (from I to ).H drog 1.5h 2 ond length g from2.4Xto3.5n rates the two nonhydrogen atoms -bond in a hydrogen bond. The strongest hydrogen bonds have a tendency to be approximately 09420 o straight,such that the hydrogen-bond donor,the hydrogen atom,and the hydrogen-bond acceptor lie along a straight line.This tendency toward lin- earity can be】 mportant I or orienting interacting molecules wi respect t the anoth Hydrogen that 3.van der Waals Interactions.The basis of a van der Waals interaction is that the distribution of electronic charge around an atom fluctuates with time.At any instant,the charge distribution is not perfectly symmetric This transient asymmetry in the electronic charge about an atom acts Distance through nteract nduce a on within its ne ng om and It come clos til the contact distance (Figure 1.10).At distances shorter than the van der Waals contact distance,very strong repulsive forces become dominant pecause the outer electron clouds of the two atoms overlap. Energies associated with van der Waals typical interactions cont 0.5 to 1 kcal P atom pa hen the s ces ol les com L er, a large number of a oms a tact tial distance. the fo the hydro action.after we examine the characteristics of water:these characteristics are essential to understanding the hydrophobic interaction. Properties of water.Water is the solvent in which most biochemical reac- tions take place,and its properties are essential to the formation of macro- o 。 trons away fron region arour

8 CHAPTER 1 Biochemistry: An Evolving Science 2. Hydrogen Bonds . These interactions are largely ionic interactions, with partial charges on nearby atoms attracting one another. Hydrogen bonds are responsible for specific base-pair formation in the DNA double helix. The hydrogen atom in a hydrogen bond is partially shared by two electronegative atoms such as nitrogen or oxygen. The hydrogen-bond donor is the group that includes both the atom to which the hydrogen atom is more tightly linked and the hydrogen atom itself, whereas the hydrogen-bond acceptor is the atom less tightly linked to the hydrogen atom (Figure 1.9). The electronegative atom to which the hydrogen atom is covalently bonded pulls electron density away from the hydrogen atom, which thus develops a partial positive charge ( d ). Thus, the hydrogen atom with a partial positive charge can interact with an atom having a partial negative charge ( d ) through an ionic interaction. Hydrogen bonds are much weaker than covalent bonds. They have energies ranging from 4 to 20 kJ mol 1 (from 1 to 5 kcal mol 1 ). Hydrogen bonds are also somewhat longer than covalent bonds; their bond lengths (measured from the hydrogen atom) range from 1.5 Å to 2.6 Å; hence, a distance ranging from 2.4 Å to 3.5 Å separates the two nonhydrogen atoms in a hydrogen bond. The strongest hydrogen bonds have a tendency to be approximately straight, such that the hydrogen-bond donor, the hydrogen atom, and the hydrogen-bond acceptor lie along a straight line. This tendency toward linearity can be important for orienting interacting molecules with respect to one another. Hydrogen-bonding interactions are responsible for many of the properties of water that make it such a special solvent, as will be described shortly. 3. van der Waals Interactions . The basis of a van der Waals interaction is that the distribution of electronic charge around an atom fluctuates with time. At any instant, the charge distribution is not perfectly symmetric. This transient asymmetry in the electronic charge about an atom acts through ionic interactions to induce a complementary asymmetry in the electron distribution within its neighboring atoms. The atom and its neighbors then attract one another. This attraction increases as two atoms come closer to each other, until they are separated by the van der Waals contact distance (Figure 1.10). At distances shorter than the van der Waals contact distance, very strong repulsive forces become dominant because the outer electron clouds of the two atoms overlap. Energies associated with van der Waals interactions are quite small; typical interactions contribute from 2 to 4 kJ mol 1 (from 0.5 to 1 kcal mol 1 ) per atom pair. When the surfaces of two large molecules come together, however, a large number of atoms are in van der Waals contact, and the net effect, summed over many atom pairs, can be substantial. We will cover the fourth noncovalent interaction, the hydrophobic interaction, after we examine the characteristics of water; these characteristics are essential to understanding the hydrophobic interaction. Properties of water. Water is the solvent in which most biochemical reactions take place, and its properties are essential to the formation of macromolecular structures and the progress of chemical reactions. Two properties of water are especially relevant: 1. Water is a polar molecule . The water molecule is bent, not linear, and so the distribution of charge is asymmetric. The oxygen nucleus draws electrons away from the two hydrogen nuclei, which leaves the region around N H N N H O O H N O H O Hydrogenbond donor Hydrogenbond acceptor + − − FIGURE 1.9 Hydrogen bonds. Hydrogen bonds are depicted by dashed green lines. The positions of the partial charges (d and d) are shown. N H O 0.9 Å 2.0 Å 180° Hydrogenbond donor Hydrogen-bond acceptor Energy Attraction Repulsion 0 van der Waals contact distance Distance FIGURE 1.10 Energy of a van der Waals interaction as two atoms approach each other. The energy is most favorable at the van der Waals contact distance. Owing to electron–electron repulsion, the energy rises rapidly as the distance between the atoms becomes shorter than the contact distance. O H H + – Electric dipole

each hydrogen atom with a net positive charge. The water molecule is thus an electrically polar structure. 2. Water is highly cohesive . Water molecules interact strongly with one another through hydrogen bonds. These interactions are apparent in the structure of ice (Figure 1.11). Networks of hydrogen bonds hold the structure together; similar interactions link molecules in liquid water and account for many of the properties of water. In the liquid state, approximately one in four of the hydrogen bonds present in ice are broken. The polar nature of water is responsible for its high dielectric constant of 80. Molecules in aqueous solution interact with water molecules through the formation of hydrogen bonds and through ionic interactions. These interactions make water a versatile solvent, able to readily dissolve many species, especially polar and charged compounds that can participate in these interactions. The hydrophobic effect. A final fundamental interaction called the hydrophobic effect is a manifestation of the properties of water. Some molecules (termed nonpolar molecules ) cannot participate in hydrogen bonding or ionic interactions. The interactions of nonpolar molecules with water molecules are not as favorable as are interactions between the water molecules themselves. The water molecules in contact with these nonpolar molecules form “cages” around them, becoming more well ordered than water molecules free in solution. However, when two such nonpolar molecules come together, some of the water molecules are released, allowing them to interact freely with bulk water (Figure 1.12). The release of water from such cages is favorable for reasons to be considered shortly. The result is that nonpolar molecules show an increased tendency to associate with one another in water compared with other, less polar and less self-associating, solvents. This tendency is called the hydrophobic effect and the associated interactions are called hydrophobic interactions . The double helix is an expression of the rules of chemistry Let us now see how these four noncovalent interactions work together in driving the association of two strands of DNA to form a double helix. First, each phosphate group in a DNA strand carries a negative charge. These negatively charged groups interact unfavorably with one another over distances. Thus, unfavorable ionic interactions take place when two strands of FIGURE 1.11 Structure of ice. Hydrogen bonds (shown as dashed green lines) are formed between water molecules to produce a highly ordered and open structure. FIGURE 1.12 The hydrophobic effect. The aggregation of nonpolar groups in water leads to the release of water molecules, initially interacting with the nonpolar surface, into bulk water. The release of water molecules into solution makes the aggregation of nonpolar groups favorable. Nonpolar molecule Nonpolar molecule Nonpolar molecule Nonpolar molecule 9 1.3 Chemical Concepts

10 DNA come together. These phosphate groups are far apart in the double helix with distances greater than 10 Å, but many such interactions take place (Figure 1.13). Thus, ionic interactions oppose the formation of the double helix. The strength of these repulsive ionic interactions is diminished by the high dielectric constant of water and the presence of ionic species such as Na or Mg 2 ions in solution. These positively charged species interact with the phosphate groups and partly neutralize their negative charges. Second, as already noted, hydrogen bonds are important in determining the formation of specific base pairs in the double helix. However, in singlestranded DNA, the hydrogen-bond donors and acceptors are exposed to solution and can form hydrogen bonds with water molecules. N H C O N H C O H O H H O H H O H H O H + When two single strands come together, these hydrogen bonds with water are broken and new hydrogen bonds between the bases are formed. Because the number of hydrogen bonds broken is the same as the number formed, these hydrogen bonds do not contribute substantially to driving the overall process of double-helix formation. However, they contribute greatly to the specificity of binding. Suppose two bases that cannot form Watson–Crick base pairs are brought together. Hydrogen bonds with water must be broken as the bases come into contact. Because the bases are not complementary in structure, not all of these bonds can be simultaneously replaced by hydrogen bonds between the bases. Thus, the formation of a double helix between noncomplementary sequences is disfavored. Third, within a double helix, the base pairs are parallel and stacked nearly on top of one another. The typical separation between the planes of adjacent base pairs is 3.4 Å, and the distances between the most closely approaching atoms are approximately 3.6 Å. This separation distance corresponds nicely to the van der Waals contact distance (Figure 1.14). Bases tend to stack even in single-stranded DNA molecules. However, the base stacking and associated van der Waals interactions are nearly optimal in a double-helical structure. Fourth, the hydrophobic effect also contributes to the favorability of base stacking. More-complete base stacking moves the nonpolar surfaces of the bases out of water into contact with each other. The principles of double-helix formation between two strands of DNA apply to many other biochemical processes. Many weak interactions contribute to the overall energetics of the process, some favorably and some unfavorably. Furthermore, surface complementarity is a key feature: when complementary surfaces meet, hydrogen-bond donors align with hydrogenbond acceptors and nonpolar surfaces come together to maximize van der Waals interactions and minimize nonpolar surface area exposed to the aqueous environment. The properties of water play a major role in determining the importance of these interactions. The laws of thermodynamics govern the behavior of biochemical systems We can look at the formation of the double helix from a different perspective by examining the laws of thermodynamics. These laws are general van der Waals contacts FIGURE 1.14 Base stacking. In the DNA double helix, adjacent base pairs are stacked nearly on top of one another, and so many atoms in each base pair are separated by their van der Waals contact distance. The central base pair is shown in dark blue and the two adjacent base pairs in light blue. Several van der Waals contacts are shown in red. FIGURE 1.13 Ionic interactions in DNA. Each unit within the double helix includes a phosphate group (the phosphorus atom being shown in purple) that bears a negative charge. The unfavorable interactions of one phosphate with several others are shown by red lines. These repulsive interactions oppose the formation of a double helix.