Ch 2 Water - Solvent for Biochemical Reactions (9th ed) PDF

Chapter 2 Water: The Solvent for Biochemical Reactions Life processes depend on the properties of Water: 2.1 Water and Polarity Page 33 Solvent Properties of Water 2.2 Hydrogen Bonds Page 38 Other Biologically Important Hydrogen Bonds The Importance of the Hydrogen Bond 2.3 Ac...

Chapter 2 Water: The Solvent for Biochemical Reactions Life processes depend on the properties of Water: 2.1 Water and Polarity Page 33 Solvent Properties of Water 2.2 Hydrogen Bonds Page 38 Other Biologically Important Hydrogen Bonds The Importance of the Hydrogen Bond 2.3 Acids, Bases, and pH Page 41 2.4 Titration Curves Page 45 Fig. 2-CO, p. 35 2.1 Water and Polarity Water is the principal component of most cells. The geometry of the water molecule and its properties as a solvent play major roles in determining the properties of living systems. The tendency of an atom to attract electrons to itself in a chemical bond (i.e., to become negative) is called electronegativity. Different elements have different electronegativity. (Table 2.1).  What is polarity? When two atoms with the same electronegativity form a bond, the electrons are shared equally between the two atoms. If atoms with differing electronegativity form a bond, the electrons are not shared equally and more of the negative charge is found closer to one of the atoms. In the O—H bonds in water, oxygen is more electronegative than hydrogen, so there is a higher probability that the bonding electrons are closer to the oxygen. This gives rise to a partial positive and negative charge. Bonds such as this are called polar bonds. In situations in which the electronegativity difference is quite small, such as in the C—H bond in methane (CH4), the sharing of electrons in the bond is very nearly equal, and the bond is essentially nonpolar. Fig. 2-1, p. 36  The bonds in a molecule may be polar, but the molecule itself can still be nonpolar because of its geometry. Carbon dioxide is an example. The two C=O bonds are polar, but because the CO2 molecule is linear, the attraction of the oxygen for the electrons in one bond is cancelled out by the equal and opposite attraction for the electrons by the oxygen on the other side of the molecule. O=C=O Water is a bent molecule with a bond angle of 104.3° and the uneven sharing of electrons in the two bonds is not cancelled out. The result is that the bonding electrons are more likely to be found at the oxygen end of the molecule than at the hydrogen end. Solvent Properties of Water Why do some chemicals dissolve in water while others do not? The polar nature of water largely determines its solvent properties for : Ionic compounds with full charges, such as potassium chloride (KCl, K+ and Cl- ; ions in solution), polar compounds with partial charges (i.e., dipoles), such as ethyl alcohol or acetone, tend to dissolve in water. The electrostatic attraction between unlike charges (the negative end of a water dipole attracts a positive ion or the positive end of another dipole). The positive end of a water molecule attracts a negative ion or the negative end of another dipole. These ion–dipole and dipole–dipole interactions are similar to the interactions between water molecules themselves in terms of the quantities of energy involved. Figure 2.2: Ionic bonds (in salts) become replaced by ion–dipole interactions. In aqueous solution, ion–dipole interactions: the cations (positively charged sodium ions) are attracted to the partial negative charges on the water. the anions (negatively charged chloride ions) are attracted to the partial positive charges on water. Water surrounding ions of this type are called hydration shells Fig. 2-2, p. 36 Fig. 2-3, p. 37  Examples of polar compounds that dissolve easily in water are small organic molecules containing one or more electronegative atoms (e.g., oxygen or nitrogen), including alcohols, amines, and carboxylic acids. The attraction between the dipoles of these molecules and the water dipoles makes them tend to dissolve. Ionic and polar substances are referred to as hydrophilic (“water-loving,” from the Greek) Hydrocarbons (compounds that contain only carbon and hydrogen) are nonpolar. It is less favorable thermodynamically for water molecules to be associated with nonpolar molecules than with other water molecules. As a result, nonpolar molecules do not dissolve in water and are referred to Fig. 2-3a, p. 37  Hydrophobic Interactions Hydrocarbons in particular tend to sequester themselves from an aqueous environment and forms a two-layer with water. The interactions between nonpolar molecules are called hydrophobic interactions or, in some cases, hydrophobic bonds. Why do oil and water mixed together separate into layers? A single molecule may have both polar (hydrophilic) and nonpolar (hydrophobic) portions. Substances of this type are called amphipathic. Example: a long-chain fatty acid having a polar carboxylic acid group: the “head” contains two oxygen atoms in addition to carbon and hydrogen, it is very polar and can form a carboxylate anion at neutral pH. a long nonpolar hydrocarbon portion: the “tail,” contains only carbon and hydrogen and is thus nonpolar (Figure 2.4). A compound such as this in the presence of water tends to form structures called micelles, in which the polar head groups are in contact with the aqueous environment and the nonpolar tails are sequestered from the water (Figure 2.5). FIGURE 2.4 Sodium palmitate, an amphiphilic molecule. the ball represents the hydrophilic polar head and the zigzag line represents the nonpolar hydrophobic hydrocarbon tail. Fig. 2-4, p. 38 FIGURE 2.5 Micelle formation by amphipathic molecules in aqueous solution. When micelles form, the ionized polar groups are in contact with the water, and the nonpolar parts of the molecule are protected from contact with the water. Fig. 2-5, p. 39  A similar process is responsible for the separation of oil and water (mixed but not dissolved), after mixing small spheres or oil droplets move to the top and coalesce into the oil layer. Interactions between nonpolar molecules themselves are very weak and depend on the attraction between short-lived temporary dipoles and the dipoles they induce. A temporary dipole can induce another dipole in a neighboring molecule in the same way that a permanent dipole does. The interaction energy is low because the association is so short-lived. It is called a van der Waals interaction (or bond). The arrangement of molecules in cells (membranes) strongly depends on the molecules’ polarity, as we saw with micelles. Summary of some important noncovalent bonds Ionic Bonds In a crystal of salt, the positive and negative ions are held together by ionic bonds. Ionic bonds and covalent bonds are the strongest bonds. polar bonds bonds in which two atoms have an unequal share in the bonding electrons nonpolar refers to a bond in which two atoms share electrons evenly dipoles molecules with positive and negative ends due to an uneven distribution of electrons in bonds salt bridge biomolecules often have ionizable groups on them and their interaction depends on the attraction of unlike charges van der Waals radius the distance between an atom’s nucleus and its effective electronic surface Hydrogen Bonds Noncovalent interaction of electrostatic origin and can be considered a special case of dipole–dipole interaction. Partial positive charge on hydrogen can interact with an unshared (nonbonding) pair of electrons (a source of negative charge) on another electronegative atom. All three atoms lie in a straight line, forming a hydrogen bond. Why does water have such interesting and unique properties? Water constitutes an optimum situation in terms of the number of hydrogen bonds that each molecule can form. Each water molecule is involved in four hydrogen bonds—as a donor in two and as an acceptor in two. Ammonia has three hydrogens to donate to a hydrogen bond but only one unshared pair of electrons, on the nitrogen. The geometric arrangement of hydrogen-bonded water molecules has important implications for the properties of water as a solvent. The result is a tetrahedral arrangement of water molecules; each molecule bind four molecules. Liquid water consists of hydrogen-bonded arrays that resemble ice crystals but in liquid water, hydrogen bonds are constantly breaking and new ones are constantly forming. An ice crystal has a more-or-less-stable arrangement of hydrogen bonds. Fig. 2-7, p. 40 numbers of Hydrogen Bonding sites (4 in water) - HF - H2 O - NH3 Hydrogen bonds (non-covalent) are much weaker than normal covalent bonds. Table 2-3, p. 41  Even this comparatively small amount of energy of H bond but it is enough to affect the properties of water especially its melting point, its boiling point, and its density relative to the density of ice. Both the melting point and the boiling point of water are significantly higher than would be predicted for a molecule of this size. The forces of attraction between the molecules of these substances are weaker than the attraction between water molecules because of the number and strength of their hydrogen bonds. Table 2-4, p. 41  Ice has a lower density than liquid water because the fully hydrogen bonded array in an ice crystal is less densely packed than that in liquid water. Most substances contract when they freeze opposite to water. Freezing and expansion of the water can kill living cells and used in laboratory procedures for cell disruption (freezing and thawing). Lakes and rivers freeze from top to bottom rather than vice versa and ice float on the surface of water. Hydrogen bonding also plays a role in the behavior of water as a solvent. If a polar solute can serve as a donor or an acceptor of hydrogen bonds: examples; alcohols, amines, carboxylic acids, and esters, as well as aldehydes and ketones, can all form hydrogen bonds with water, so they are soluble in water. Fig. 2-8, p. 41 Fig. 2-9, p. 42  Hydrogen bonds have a vital involvement in stabilizing the three- dimensional structures of biologically important molecules, including DNA, RNA, and proteins. The hydrogen bonds between complementary bases in the double- helical structure of DNA. Hydrogen bonding in proteins gives rise to two important structures, the a-helix and b-pleated sheet conformations. p. 43  2.3 Acids, Bases, and pH (page 47) 2.4 Titration Curves (page51) 2.5 Buffers (page 53) An acid is a molecule that acts as a proton (H+; hydrogen ion) donor. A base is a proton acceptor. How easy acids or bases lose or gain protons depends on the chemical nature of the compounds. The degree of dissociation of acids in water, ranges from essentially complete dissociation for a strong acid to little dissociation for a very weak acid, and any intermediate value. The numerical measure of acid strength (which is the amount of H+ released in water) called acid dissociation constant (Ka). Ka has a fixed numerical value at a given temperature. The greater the Ka, the stronger the acid. Water acts as a base as well as the solvent. Table 2-5, p. 43 Table 2-6, p. 46  What is pH? The acid–base properties of water (solvent) play an important part in biological processes. The self-dissociation of water to: hydroxide ion + hydrogen ion This relationship is derived for pure water and is valid for any aqueous solution (neutral, acidic, or basic). The pH is measurement of this balance. One pH unit change = a tenfold difference in H+ concentration. FIGURE 2.12 pH versus enzymatic activity. Pepsin, trypsin, and lysozyme all have steep pH optimum curves. Pepsin has maximum activity under very acidic conditions (a digestive enzyme found in the stomach). Lysozyme has its maximum activity near pH 5. Trypsin is most active near pH 6. Fig. 2-12, p. 47  If we dissolved a weak acid (has a known Ka) in an aqueous solution, the amount of acid and its conjugate base depends on the pH of that solution. The equation connects the Ka with the pH of a solution or [H+] is called Henderson–Hasselbalch equation. It is useful in predicting the properties of buffer solutions used to control the pH of reaction mixtures. The equilibrium indicates that increasing [A-] means less [HA] and vice versa. EXAMPLE: [HA] = [A-]; the ratio [A]/[HA] = 1; the logarithm of 1= 0; the pH of that solution equals the pK value of the weak acid. a 2.4 Titration Curves A titration is an experiment in which measured amounts of base are added to a measured amount of acid and follow the course with a pH meter. The point in the titration at which the acid is exactly neutralized is called the equivalence point. If the pH is monitored as base is added to a sample of acetic acid in the course of a titration, an inflection point in the titration curve is reached when the pH equals the pKa of acetic acid (= 4.76). A region near the pKa; the curve is relatively flat (the pH changes very little as base is added in this region. Based on Henderson–Hasselbalch equation, a pH value equal to the pKa of the acid (4.76); [weak acid] = [conjugate base] [acetic acid] = [acetate ion] The inflection point occurs when 0.5 mol of base has been added for each mole of acid present. The acid (acetic acid) is progressively converted to its conjugate base (acetate ion) as more NaOH is added and the titration proceeds. At the end of titration; all the acetic acid has been converted to acetate ion. The form of this Figure represents the behavior of any monoprotic weak acid. The value of the pKa for each individual acid determines the pH values at the inflection point and at the equivalence point.  Acids are categorized into three groups: monoprotic acids (release one hydrogen ion, have a single Ka and pKa). diprotic acids (release two hydrogen ions, have two Ka values and two pKa values). polyprotic acids (release more than two hydrogen ions). e.g. citric acid and phosphoric acid (can release three hydrogen ions and have three Ka values and three pKa values). (Amino acids, behave as diprotic and polyprotic acids). When the pH of a solution equal the pKa of an acid, the protonated and deprotonated forms are equal. pH = pKa protonated = deprotonated When the pH of a solution is less than the pKa of an acid, the protonated form predominates. pH < pKa H+ on, substance protonated When the pH of a solution is greater than the pKa of an acid, the deprotonated (conjugate base) form predominates. pH > pKa H + off, substance deprotonated 2.5 Buffers A buffer is something that resists change. pH buffer solution tends to resist change in pH when small to moderate amounts of a strong acid or strong base are added. A buffer solution consists of a mixture of a weak acid and its conjugate base. How do buffers work? Triprotic Phosphoric Acid Release more 3 proterons and have three Ka values and three pKa values). FIGURE 2.15 The relationship between the titration curve and buffering action in H2PO4 Fig. 2-15, p. 53 Fig. 2-15a, p. 53 Fig. 2-15b, p. 53  A buffer solution can maintain the pH at a relatively constant value because of the presence of appreciable amounts of both the acid and its conjugate base. This condition is met at pH values at or near the pKa of the acid. If OH is added, an appreciable amount of the acid form of the buffer is present in solution to react with the added base. If H is added, an appreciable amount of the basic form of the buffer also is present to react with the added acid. Table 2-7, p. 54  At pH values below the pKa, the acid form predominates, At pH values above the pKa, the basic form predominates. The plateau region in a titration curve, where the pH does not change rapidly, covers a pH range extending approximately one pH unit on each side of the pK. Thus, the buffer is effective within a range of a about two pH units. Fig. 2-16, p. 54  In many biochemical studies, a strict pH range must be maintained in order for the experiment to be successful. We can select an appropriate buffer from pKa e.g. if we needed the pH = 7.2, we might select the H2PO4/HPO4 pair to be our buffer. e.g. If we wanted a pH near 9.0, we would look at tables of buffers to find one with a pKa close to nine. How do we make buffers in the laboratory? 1. In theory, we can use the HendersonHasselbalch equation and do calculations concerning ratios of conjugate base form to conjugate acid form. Adding predetermined amounts of the conjugate base form (A) to the acid form (HA). 2. We could start with one and create the other. HA and A are interconverted by adding strong acid or strong base. To make a buffer, we could start with the HA form and add NaOH until the pH is correct, as determined by a pH meter. Or could also start with A and add HCl. Fig. 2-14, p. 50 Table 2-8, p. 56 Naturally occurring pH buffers present in living organisms Buffer systems in living organisms are based on many types of compounds. pH in most organisms around 7 In blood, phosphate ion levels are inadequate for buffering, and a different system operates. The buffering system in blood is based on the dissociation of carbonic acid (H 2CO3); pKa=6.37 The pH of human blood, 7.4, is near the end of the buffering range of this system. Carbon dioxide can dissolve in water and in water-based fluids, such as blood. The dissolved carbon dioxide forms carbonic acid, which, in turn, reacts to produce bicarbonate ion:  At the pH of blood, which is about one unit higher than the pKa of carbonic acid, most of the dissolved CO 2 is present as HCO3. The CO2 being transported to the lungs to be expired takes the form of bicarbonate ion. A direct relationship exists between the pH of the blood and the pressure of carbon dioxide gas in the lungs. Hemoglobin, the oxygen-carrying protein in the blood, also play role in buffering blood due to histidine amino acid. Biochemical laboratory buffers that have come into wide use more recently are zwitterions, which are compounds that have both a positive charge and a negative charge. Zwitterions are usually considered less likely to interfere with biochemical reactions than some of the earlier buffers Most living systems operate at pH levels close to 7. The pKa values of important functional groups, such as the carboxyl and amino groups, are well above or well below this value. As a result, under physiological conditions, many important biomolecules exist as charged specie. p. 57