1. Water & pH

2. Water The predominant chemical component of living organisms
The ability to solvate a wide range of organic and inorganic molecules
Dipolar structure
a high dielectric constant
Reflecting the number of dipoles
Capacity for forming hydrogen bonds

3. Water Influences the structure of biomolecule
Water is a reactant or product in many metabolic reactions

4. Water
Water has a slight propensity to dissociate into hydroxide ions and protons
The acidity of aqueous solutions is generally reported using the logarithmic pH scale
Buffers normally maintain the pH of extracellular fluid between 7.35 and 7.45.
Suspected disturbances of acid-base balance are verified by measuring the pH of arterial blood and the CO2 content of venous blood

5. WATER IS AN IDEAL BIOLOGIC SOLVENT
Water Molecules Form Dipoles

6. FIGURE 2-1a Structure of the water molecule
FIGURE 2-1a Structure of the water molecule. (a) The dipolar nature of the H2O molecule is shown in a ball-and-stick model; the dashed lines represent the nonbonding orbitals. There is a nearly tetrahedral arrangement of the outer-shell electron pairs around the oxygen atom; the two hydrogen atoms have localized partial positive charges (δ+) and the oxygen atom has a partial negative charge (δ–).

7. WATER IS AN IDEAL BIOLOGIC SOLVENT
the strength of interaction F between oppositely charged particles is inversely proportionate to the dielectric constant ε of the surrounding medium.
Water therefore greatly decreases the force of attraction between charged and polar species relative to water-free environments with lower dielectric constants.

8. Water Molecules Form Hydrogen Bonds

9. FIGURE 2-1b Structure of the water molecule
FIGURE 2-1b Structure of the water molecule. (b) Two H2O molecules joined by a hydrogen bond (designated here, and throughout this book, by three blue lines) between the oxygen atom of the upper molecule and a hydrogen atom of the lower one. Hydrogen bonds are longer and weaker than covalent O—H bonds.

10. FIGURE 2-3 Common hydrogen bonds in biological systems
FIGURE 2-3 Common hydrogen bonds in biological systems. The hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom.

11. FIGURE 2-4 Some biologically important hydrogen bonds.

12. Biomolecules,have functional groups

15. Water as solvent
FIGURE 2-6 Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the C– and Na+ ions. The ionic charges are partially neutralized, and the electrostatic attractions necessary for lattice formation are weakened.

16. FIGURE 2-7b (part 3) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 3) Amphipathic compounds in aqueous solution. (b) By clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle.

17. FIGURE 2-5 Directionality of the hydrogen bond
FIGURE 2-5 Directionality of the hydrogen bond. The attraction between the partial electric charges (see Figure 2-1) is greatest when the three atoms involved in the bond (in this case O, H, and O) lie in a straight line. When the hydrogen-bonded moieties are structurally constrained (when they are parts of a single protein molecule, for example), this ideal geometry may not be possible and the resulting hydrogen bond is weaker.

18. Hydrophilic Hydrophobic
The forces that hold the nonpolar regions of the molecules together are called hydrophobic interactions.

19. Most biomolecules are amphipathic;
Proteins
Phospholipid bilayer
Biomolecules fold to position polar & charged groups on their surfaces
It minimizes energetically unfavorable contact between water and hydrophobic groups

20. WATER IS AN EXCELLENT NUCLEOPHILE
Nucleophilic attack by water generally results in the cleavage of the amide, glycoside, or ester bonds
Hydrolysis

21. Water as a product
when monomer units are joined together to form biopolymers
such as proteins or glycogen

22. FIGURE 2-12 Effect of extracellular osmolarity on water movement across a plasma membrane. When a cell in osmotic balance with its surrounding medium—that is, a cell in (a) an isotonic medium—is transferred into (b) a hypertonic solution or (c) a hypotonic solution, water moves across the plasma membrane in the direction that tends to equalize osmolarity outside and inside the cell.

23. FIGURE 2-11 Osmosis and the measurement of osmotic pressure
FIGURE 2-11 Osmosis and the measurement of osmotic pressure. (a) The initial state. The tube contains an aqueous solution, the beaker contains pure water, and the semipermeable membrane allows the passage of water but not solute. Water flows from the beaker into the tube to equalize its concentration across the membrane. (b) The final state. Water has moved into the solution of the nonpermeant compound, diluting it and raising the column of water within the tube. At equilibrium, the force of gravity operating on the solution in the tube exactly balances the tendency of water to move into the tube, where its concentration is lower. (c) Osmotic pressure (Π) is measured as the force that must be applied to return the solution in the tube to the level of that in the beaker. This force is proportional to the height, h, of the column in (b).

24. The effect of solutes on osmolarity depends on the number of dissolved particles, not their mass
The high concentration of albumin and other proteins in blood plasma contributes to its osmolarity.
Cells also actively pump out ions such as Na into the interstitial fluid to stay in osmotic balance with their surroundings.

25. Regulation of water balance
Hypothalamic mechanisms that control thirst
Antidiuretic hormone (ADH)
Retention or excretion of water
Kidneys, and on evaporative loss
Nephrogenic diabetes
Inability to concentrate urine
Unresponsiveness of renal tubular osmoreceptors to ADH.

26. Acidity of aqueous solutions (pH)
pH = −log [H+ ]
Low pH values
Correspond to high concentrations of H+
High pH values
Correspond to low concentrations of H+
Acids
are proton donors
Bases
are proton acceptors

27. K w=[H+][OH−]=10−14
pH+pOH=14
pH = −log [H+ ]

28. pH Strong acids Weak acids dissociate only partially
Have larger dissociation constants
Weak acids dissociate only partially
Strong bases (eg, KOH or NaOH)
the strong base KOH is completely dissociated in solution and that the concentration of OH ions is thus equal to that of the KOH
Weak bases (eg, Ca[OH]2)
Many biochemicals are weak acids

29. Many biochemicals possess functional groups that are weak acids or bases
We express the relative strengths of weak acids and bases in terms of their dissociation constants (Ka)
expressing the extent of ionization of water in quantitative terms.
The equilibrium constant is fixed and characteristic for any given chemical reaction at a specified temperature.

30. The tendency of any acid (HA) to lose a proton and form its conjugate base (A-) is defined by the equilibrium constant

31. pKa is used to express the relative strengths of both acids and bases.
polyproteic compounds
Containing more than one dissociable proton, a numerical subscript is assigned to each in order of relative acidity
the pKa is the pH at which the concentration of the acid (R-NH3+) equals that of the base (R-NH2).

34. The pKa Values Depend on
pKa of an acid group is the pH at which the protonated and unprotonated species are present at equal concentrations
The pKa Values Depend on
Molecular Structure
The presence of adjacent negative charge
Decreases with distance
Properties of the Medium

35. Buffer(s)
Solutions of weak acids or bases and their conjugates exhibit buffering
Maximum buffering capacity
Most effectively in the pH range pKa ± 1.0 pH unit.
Physiologic buffers
The value of pKa relative to the desired pH is the major determinant of which buffer is selected.

36. The pH of an aqueous solution can be approximately measured using
indicator dyes including
Litmus, phenolphthalein, and phenol red, which undergo color changes as a proton dissociates from the dye molecule
a glass electrode that is selectively sensitive to H+ concentration

37. The pH of some aqueous fluids
FIGURE 2-14 The pH of some aqueous fluids.

38. Disturbances of acid-base balance
Measuring the pH of arterial blood
The CO2 content of venous blood
Acidosis (blood pH < 7.35)
Causes
include diabetic ketosis
Lactic acidosis
Alkalosis (pH > 7.45)
Vomiting of acidic gastric contents

40. Weak Interactions Are Crucial to Macromolecular Structure and Function
Macromolecules such as proteins, DNA, and RNA contain so many sites of potential hydrogen bonding or ionic, van der Waals, or hydrophobic interactions
the cumulative effect of the many small binding forces can be enormous.
For macromolecules, the most stable (that is, the native) structure is usually that in which weak-bonding possibilities are maximized

41. The folding of a single polypeptide or polynucleotide chain into its three-dimensional shape
The binding of an antigen to a specific antibody
Formation of an enzyme-substrate complex
The binding of a hormone or a neurotransmitter to its cellular receptor protein

42. FIGURE 2-8 Release of ordered water favors formation of an enzyme-substrate complex. While separate, both enzyme and substrate force neighboring water molecules into an ordered shell. Binding of substrate to enzyme releases some of the ordered water, and the resulting increase in entropy provides a thermodynamic push toward formation of the enzyme-substrate complex (see p. 192).

43. Titration curve for an acid of the type
HA. The heavy dot in the center of the curve indicates
the pKa 5.0.

44. FIGURE 2-16 The titration curve of acetic acid
FIGURE 2-16 The titration curve of acetic acid. After addition of each increment of NaOH to the acetic acid solution, the pH of the mixture is measured. This value is plotted against the amount of NaOH added, expressed as a fraction of the total NaOH required to convert all the acetic acid (CH3COOH) to its deprotonated form, acetate (CH3COO–). The points so obtained yield the titration curve. Shown in the boxes are the predominant ionic forms at the points designated. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the pKa. The shaded zone is the useful region of buffering power, generally between 10% and 90% titration of the weak acid.

45. H+(H2O)n Water Is a Weak Electrolyte
Protons that dissociate interact with oxygens of other water molecules to form clusters of water molecules.
H+(H2O)n

47. The dissociation of an acid increases with increasing temperatures.
Keq will be a
Small number if the degree of dissociation of a substance is small.
Large if the degree of dissociation is large

48. An extremely small number of water molecules actually dissociate
At 25°C the value of Keq for dissociation of water is about 1.8 x 10-16

52. The relationship between pK' and Keq is an inverse one
The smaller the Keq the larger the pK’

54. The Relationship between pH and Concentrations of Conjugate Acid and Base
If [base]/[acid] is 1: 1, pH equals the pK’

55. Must be taken into account for a buffer
(best)Buffering range
Between 1 pH unit below and 1 pH unit above pK‘.
Buffering capacity
Depends on concentrations of conjugate acid and base

57. To monitor the acid-base parameters of a patient's blood.
Values of interest to a clinician
pH,HC03 - and CO2 concentrations.
Normal values
pH = 7.40, [HC03 -] =24.0 mM, [C02] =1.20 mM.