H. Lodish - Molecular Cell Biology (5ed, Freeman, 2003) (796244), страница 15
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The sizes and directions of thedipole moments of each of the bonds determine the net dipolemoment of the molecule.Covalent Bonds Are Much Stronger and MoreStable Than Noncovalent InteractionsCovalent bonds are very stable because the energies requiredto break them are much greater than the thermal energyavailable at room temperature (25 ºC) or body temperature(37 ºC). For example, the thermal energy at 25 ºC is approximately 0.6 kilocalorie per mole (kcal/mol), whereas theenergy required to break the carbon-carbon single bond(COC) in ethane is about 140 times larger (Figure 2-4). Consequently at room temperature (25 ºC), fewer than 1 in 1012ethane molecules is broken into a pair of ·CH3 radicals, eachcontaining an unpaired, nonbonding electron.Covalent single bonds in biological molecules have energies similar to that of the COC bond in ethane.
Becausemore electrons are shared between atoms in double bonds,they require more energy to break than single bonds. For instance, it takes 84 kcal/mol to break a single COO bond, but170 kcal/mol to break a CUO double bond. The most common double bonds in biological molecules are CUO, CUN,CUC, and PUO.The energy required to break noncovalent interactionsis only 1–5 kcal/mol, much less than the bond energies ofcovalent bonds (see Figure 2-4).
Indeed, noncovalent interactions are weak enough that they are constantly beingHIn the resonance hybrid on the right, one of the electronsfrom the PUO double bond has accumulated around the Oatom, giving it a negative charge and leaving the P atom witha positive charge. These charges are important in noncovalent interactions.Noncovalent interactions▲FIGURE 2-4 Relative energies ofcovalent bonds and noncovalentinteractions.
Bond energies are determinedas the energy required to break a particulartype of linkage. Covalent bonds are one totwo powers of 10 stronger than noncovalentinteractions. The latter are somewhat greaterthan the thermal energy of the environmentat normal room temperature (25 ˚C). Manybiological processes are coupled to theenergy released during hydrolysis of aphosphoanhydride bond in ATP.OCovalent bondsElectrostaticvan derWaalsHydrogenbondsThermalenergy0.24 × 100Hydrolysis of ATPphosphoanhydride bond0.24 × 1010.24 × 102C−CC=C0.24 × 103kcal/mol2.1 • Atomic Bonds and Molecular Interactions33formed and broken at room temperature.
Although theseinteractions are weak and have a transient existence atphysiological temperatures (25–37 ºC), multiple noncovalent interactions can act together to produce highly stableand specific associations between different parts of a largemolecule or between different macromolecules. We first review the four main types of noncovalent interactions andthen consider their role in the binding of biomolecules toone another and to other molecules.The relative strength of the interaction between two ions,A and C, depends on the concentration of other ions in asolution.
The higher the concentration of other ions (e.g., Naand Cl), the more opportunities A and C have to interactionically with these other ions, and thus the lower the energyrequired to break the interaction between A and C. As aresult, increasing the concentrations of salts such as NaCl ina solution of biological molecules can weaken and even disrupt the ionic interactions holding the biomolecules together.Ionic Interactions Are AttractionsBetween Oppositely Charged IonsHydrogen Bonds Determine Water Solubilityof Uncharged MoleculesIonic interactions result from the attraction of a positivelycharged ion—a cation—for a negatively charged ion—ananion. In sodium chloride (NaCl), for example, the bondingelectron contributed by the sodium atom is completely transferred to the chlorine atom.
Unlike covalent bonds, ionicinteractions do not have fixed or specific geometricorientations, because the electrostatic field around an ion—its attraction for an opposite charge—is uniform in alldirections.In aqueous solutions, simple ions of biological significance, such as Na, K, Ca2, Mg2, and Cl, do not existas free, isolated entities.
Instead, each is hydrated, surrounded by a stable shell of water molecules, which are heldin place by ionic interactions between the central ion and theoppositely charged end of the water dipole (Figure 2-5).Most ionic compounds dissolve readily in water because theenergy of hydration, the energy released when ions tightlybind water molecules, is greater than the lattice energy thatstabilizes the crystal structure. Parts or all of the aqueous hydration shell must be removed from ions when they directlyinteract with proteins.
For example, water of hydration islost when ions pass through protein pores in the cell membrane during nerve conduction (Chapter 7).A hydrogen bond is the interaction of a partially positivelycharged hydrogen atom in a molecular dipole (e.g., water)with unpaired electrons from another atom, either in thesame (intramolecular) or in a different (intermolecular) molecule. Normally, a hydrogen atom forms a covalent bondwith only one other atom. However, a hydrogen atom covalently bonded to an electronegative donor atom D may forman additional weak association, the hydrogen bond, with anacceptor atom A, which must have a nonbonding pair ofelectrons available for the interaction:Mg2HOH▲ FIGURE 2-5 Electrostatic interaction between water anda magnesium ion (Mg2).
Water molecules are held in placeby electrostatic interactions between the two positive chargeson the ion and the partial negative charge on the oxygen of eachwater molecule. In aqueous solutions, all ions are surrounded bya similar hydration shell.DH ADHAHydrogen bondThe length of the covalent DOH bond is a bit longer than itwould be if there were no hydrogen bond, because the acceptor “pulls” the hydrogen away from the donor. An important feature of all hydrogen bonds is directionality. In thestrongest hydrogen bonds, the donor atom, the hydrogenatom, and the acceptor atom all lie in a straight line. Nonlinear hydrogen bonds are weaker than linear ones; still, multiple nonlinear hydrogen bonds help to stabilize thethree-dimensional structures of many proteins.Hydrogen bonds are both longer and weaker than covalent bonds between the same atoms.
In water, for example,the distance between the nuclei of the hydrogen and oxygenatoms of adjacent, hydrogen-bonded molecules is about 0.27nm, about twice the length of the covalent OOH bondswithin a single water molecule (Figure 2-6a). The strengthof a hydrogen bond between water molecules (approximately 5 kcal/mol) is much weaker than a covalent OOHbond (roughly 110 kcal/mol), although it is greater than thatfor many other hydrogen bonds in biological molecules (1–2kcal/mol).
The extensive hydrogen bonding between watermolecules accounts for many of the key properties of thiscompound, including its unusually high melting and boilingpoints and its ability to interact with many other molecules.The solubility of uncharged substances in an aqueous environment depends largely on their ability to form hydrogenbonds with water. For instance, the hydroxyl group (OOH)in methanol (CH3OH) and the amino group (ONH2) inmethylamine (CH3NH2) can form several hydrogen bondswith water, enabling these molecules to dissolve in water to34CHAPTER 2 • Chemical Foundations(b)(a)(c)OHHHOHHHHOHHHOOWater-waterHHOCNHOHNHCH3Methanol-water▲ FIGURE 2-6 Hydrogen bonding of water with itselfand with other compounds. Each pair of nonbonding outerelectrons in an oxygen or nitrogen atom can accept a hydrogenatom in a hydrogen bond.
The hydroxyl and the amino groups canalso form hydrogen bonds with water. (a) In liquid water, eachwater molecule apparently forms transient hydrogen bonds withhigh concentrations (Figure 2-6b). In general, molecules withpolar bonds that easily form hydrogen bonds with water canreadily dissolve in water; that is, they are hydrophilic. Manybiological molecules contain, in addition to hydroxyl andamino groups, peptide and ester groups, which form hydrogen bonds with water (Figure 2-6c).
X-ray crystallographycombined with computational analysis permits an accuratedepiction of the distribution of electrons in covalent bondsand the outermost unbonded electrons of atoms, as illustrated in Figure 2-7. These unbonded electrons can form hydrogen bonds with donor hydrogens.CH3HHOHOHHOHOOOHHHHHOHCPeptide group−waterMethylamine-waterOHOEster group−waterseveral others, creating a dynamic network of hydrogen-bondedmolecules. (b) Water also can form hydrogen bonds withmethanol and methylamine, accounting for the high solubility ofthese compounds. (c) The peptide group and ester group, whichare present in many biomolecules, commonly participate inhydrogen bonds with water or polar groups in other molecules.HNCOVan der Waals Interactions AreCaused by Transient DipolesWhen any two atoms approach each other closely, they create a weak, nonspecific attractive force called a van derWaals interaction.
These nonspecific interactions result fromthe momentary random fluctuations in the distribution of theelectrons of any atom, which give rise to a transient unequaldistribution of electrons. If two noncovalently bonded atomsare close enough together, electrons of one atom will perturbthe electrons of the other. This perturbation generates a transient dipole in the second atom, and the two dipoles will attract each other weakly (Figure 2-8). Similarly, a polarcovalent bond in one molecule will attract an oppositely oriented dipole in another.Van der Waals interactions, involving either transientlyinduced or permanent electric dipoles, occur in all types ofmolecules, both polar and nonpolar. In particular, van derWaals interactions are responsible for the cohesion betweenmolecules of nonpolar liquids and solids, such as heptane,CH3O(CH2)5OCH3, that cannot form hydrogen bonds orionic interactions with other molecules.
The strength of vander Waals interactions decreases rapidly with increasing distance; thus these noncovalent bonds can form only whenCα▲ FIGURE 2-7 Distribution of bonding and outer nonbonding electrons in the peptide group. Shown here is oneamino acid within a protein called crambin. The black linesdiagrammatically represent the covalent bonds between atoms.The red (negative) and blue (positive) lines represent contours ofcharge.
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