L. David Nelson and Michael M. Cox
Lehninger Principles of Biochemistry
Fourth Edition
Chapter 2: Water
Copyright 2004 by W. H. Freeman & Company
Introduction
Water is the the most abundant substance in living system, making up 70% or more of the weight of most organism
2
Weak Interactions in Aqueous System
Hydrogen bonds (H-bonds) between water molecules provide the cohesive forces that make a liquid at room temperature and that favor the extreme ordering of molecules that is typical of crystalline water (ice). Polar biomolecules dissolves readily in water because they can replace waterwater interaction with more energetically favorable water-solute interaction. In contrast, nonpolar biomolecules interfere with water-water interactions but are unable to form water-solute interactions consequently, nonpolar molecules are poorly soluble in water.
In aqueous solutions, nonpolar molecules tend to cluster together.
Very important interactions in aqueous systems include;
H-bonds Ionic bonds Hydrophobic interactions Van de Waals (individually weak, but collectively very influential)
i.
H-bonding Water has very unusual melting point, boiling point and heat of vaporization than most solvents (TABLE 2-1). WHY?
Because of the h-bonding between the oxygen atom of one water molecule and the hydrogen of another (Figure 2-1), extra energy is require to destabilize these bonds.
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Weak Interactions in Aqueous System
i.
H-bonding (cont.) In ice, each water molecule is fixed in space and forms h-bonds with four other water molecules to yield a regular lattice structure (Fig. 2-2).
Breakage of a sufficient number of H-bonds to destabilize the crystal lattice of ice requires much thermal energy, which accounts for the relatively high melting point of water.
ii.
H-bonding between and polar solvents H-bonding readily from between an electronegative atom (H, acceptor, usually O or N with lone pairs of electrons) and a hydrogen atom covalently bonded to another electronegative atom (hydrogen donor) in the same or another molecule (Fig. 2-3). Hydrogen covalently bonded to C atoms do not participate in H-bonding.
B.P of butanol is 117C whereas butane is only 0.5 C. WHY?
Butanol has a polar hydroxyl group thus can form intramolecular H-bonds.
Alcohols, aldehydes, ketones, and compounds containing N-H bonds with water (Fig. 2-4) and tends to be soluble. Hydrogen bonds are strongest when the bonded molecules are oriented to 6 maximize electrostatic interaction (Fig. 2-5).
Weak Interactions in Aqueous System
iii. Water interacts electrostatically with charged solutes Compounds that dissolved in water are hydrophilic. Nonpolar compounds, which are poorly soluble in water are said to be hydrophobic (Table 2-2). Water dissolves NaCl by hydrating and stabilizing the Na + and Cl- ions, weakening the electrostatic interaction between them and this counteracting their tendency to associate in a crystalline lattice (Fig. 2-6).
The same applies to all biomolecules with charged groups.
F = Q1Q2/ r2 F = Force, Q = magnitude of charge, r = distant btw charges, = dielectric constant. For water at 25C, is78.5 and nonpolar benzene is 4.6.
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Weak Interactions in Aqueous System iv. Nonpolar compounds force energetically unfavorable chnages in the structure of water. Amphipathic compound contain regions that are polar (or charged) and regions that are nonpolar. When an amphipathic compound is mixed with water, the hydrophilic region dissolves in water, but the hydrophobic regions avoid water (Fig. 27a).
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Weak Interactions in Aqueous System
iv.
v.
Nonpolar compounds force energetically unfavorable chnages in the structure of water (cont). The hydrophobic regions of the molecule cluster together to present the smallest hydrophobic area to the aqueous solvent, and these regions are arranged to maximize their interactions with solvent (hydrophobic interactions) [Fig. 2-7b]. The structure formed with amphipathic compound in water are called micelles. Many biomolecules are amphipathic. Release of ordered water when substrates bind enzymes is the driving forces during substrate-enzyme interaction (Fig. 2-8) Van de Waals interactions are weak interatomic attraction. Weak interactions formed as random variations in positions of the electrons around a nucleus creates a transient dipole, which induces a transient, opposite electric dipole in the nearby atom. Each atom has a characteristic van de Waals radius, a measure of how close that atom will allow another to approach. Table 2-4 summarizes weak interaction in biomolecules.
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Ionization of Water, Weak Acids and Weak Bases
i. Ionization of water and equilibrium constant (Keq). Water molecules have a tendency to undergo reversible ionization to yield a hydrogen ion (proton) and hydroxide ion.
H2O = H+ + OHH+ formed in water are immediately hydrated to hydronium ions (H3O-) Keq = [H+][OH-]/[H2O] In pure water at 25C the concentration 55.5 M [H+] = [OH-] = 1.0 10-7 M Keq = [H+][OH-]/55.5 (55.5M)(Keq) = [H+][OH-] = Kw Ionization product of water Keq of pure water = 1.8 10-16M at 25C Therefore Kw = [H+][OH-] = (55.5M)(1.8 10-16M ) = 1.0 10-14M2 At neutral pH: Kw = [H+][OH-] = [H+]2 Solving for [H+] = (Kw) = (1.0 10-14M2) [H+] = [OH-] = 1.0 10-7 M
ii.
The Scale pH Designates the [H+] and [OH-] concentrations. The ion product of water, Kw is the basis for pH (Table 2-6).
pH = log 1/ [H+] = -log [H+] At 25C pH = log 1/ [1.0 10-7 ]= -log [1.0 10-7] = log 1.0 + log 107 = 0 + 7 = 7
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Ionization of Water, Weak Acids and Weak Bases ii. The pH Scale Designates the [H+] and [OH-] concentrations (cont). pH of common aqueous fluid are shown Fig.2-15. Measurement of pH is one of the most important and frequently used procedure in biochemistry.
pH affects the structure and activity of enzymes. pH measurement of blood and urine are commonly used in diagnosis. Example.
The pH of the blood plasma of severely diabetic people, for example, is often below the normal value of 7.4; (acidosis) or higher than normal alkalosis.
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Ionization of Water, Weak Acids and Weak Bases
iii. Weak acids and bases have characteristic constants Strong acids (HCl, sulfuric acid and nitric acid) and strong bases (NaOH &KOH) are completely ionized in dilute aqueous solutions. Weak acids and weak bases do not completely ionize when dissolved in water. Acid is a proton donor and base is proton acceptor. Proton donor and its corresponding proton acceptor make up a conjugate acid-base pair (Fig. 2-16). Example: acetic acid (CH3COOH), a proton donor and acetate anion, proton acceptor.
CH3COOH = H+ + CH3COO27
Ionization of Water, Weak Acids and Weak Bases iii. Weak acids and bases have characteristic constants (Cont.) General equation
HA = H+ + A Keq = [H+][A-]/[HA] = Ka = dissociation constant Strong acids, such as phosphoric acid and carbonic acids have larger dissociation constant. Weak acids, such as monohydrogen phosphate (HPO42) have smaller dissociation constants. pKa is analogues to pH pKa = log 1/ [Ka]= -log [Ka]
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Ionization of Water, Weak Acids and Weak Bases
iv. Titration curves reveal the pKa of weak acids Titration is used to determine the amount of acid in a given solution. A measured volume of the acid is titrated with a solution of a strong base, usually NaOH, of known concentration. NaOH is added in small increments until the acid is consumed (neutralized). The concentration of the acid in the original solution can be calculated from the volume and concentration of NaOH added. Plot of pH vs volume of NaOH added (titration curve) reveal the pKa of the weak acid.
30
Ionization of Water, Weak Acids and Weak Bases
iv.
Titration curves reveal the pKa of weak acids (cont.) Consider of the titration of 0.1 M solution of acetic acid (HAc) with 0.1 NaOH at 25C (Fig. 2-17). H2O = H+ + OH- ---------------- eq.1
HAc = H+ + Ac- ------------------ eq. 2
At equilibrium
Kw = [H+ ][OH-] = 1.0 10-7 M Ka = [H+][Ac-]/[HAc] = 1.74 10-5 M
As NaOH is gradually introduced, the added OH- combines with free H+ in solution to form H2O (equation 1). As free H+ is removed, HAc dissociate to satisfy its equilibrium (eq. 2). The net result is that as the titration proceeds more HAc is ionized. At the mid-point (0.5 of NaOH equivalent, Fig. 2-17), one-half of HAc has undergone dissociation i.e., [HAc] = [Ac-]. At the mid-point pH of equimolar solution of HAc and acetate is exactly equal to pKa of acetic acid (pKa = 4.76). As the titration is continued the remaining undissociated acetic acid is gradually converted into acetate, the end-point is at pH 7.0.
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Compares the titration curves of three weak acids with different Ka Acetic acid is the strongest and ammonium ion is the weakest.
Note: Weak acid and its anion- conjugate acid-base pair can act as a buffer
Buffering against pH Changes in Biological System
Buffer are aqueous system that tend to resist changes in pH when small amounts of acids (H+) or base (OH-) are added. Buffer system consists of a weak acid (proton donor) and its conjugate base (proton acceptor). Buffering results from two reversible reactions equilibria occuring in a solution of nearly equal concentrations of a proton donor and its conjugate (Fig. 2-19).
Whenever H+ or OH- is added, the small change in ratio of the relative concentrations of the weak acid and its anion small change in pH. The decrease in concentration of one component of the system is balanced exactly by an increase in other.
Each conjugate acid-base pair has a characteristic pH zone in which it is an effective buffer (Fig 2-18) i. pH, pKa and Buffer concentration Henderson-Hesselbalch equation.
Ka = [H+][A-]/[HA] Solve of [H+] = Ka[HA]/[A-] -log [H+] = -log Ka -log [HA]/[A-] pH = pKa + log[A-]/[HA] pH = pKa + log[Proton acceptor]/Proton donor] At Mid-point [A-] =[HA]; pH =pKa
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Buffering against pH Changes in Biological System
ii. Weak acids or bases buffer cells and tissues against pH changes Two important biological buffers are;
Phosphate buffer systems (Ranges between 5.9 7.9, effective in biological fluids, mammals).
H2PO4- = H+ + HPO42-
Bicarbonate buffer systems (Blood Plasma).
H2CO3 = H+ + HCO3K1 = [H+][HCO3-]/[H2CO3] -------equation-1 Because carbonic acid is formed from dissolved (d) CO2 and water CO2(d) + H2O = H2CO3 K2 =[H2CO3]/[CO2(d) ][H2O]----------equation-2 CO2(g) = CO2(d)
K3 =[CO2(d)]/[CO2(g)]----------equation-3
Note: pH depend on the concentrations H2CO3, HCO3-, CO2(d) and pCO2(g). Effect of pH on enzyme activity is shown in Fig. 2-21. 36