1. Water: The solvent for Biochemical Reactions
Chapter 2.
Water: The solvent for Biochemical Reactions

2. Water and Polarity What is polarity?
Chapter 2
Water and Polarity
What is polarity?
- Electronegativity: The tendency of an atom to attract electrons to itself in a chemical bond
- The difference in electronegativity causes a partial positive and negative charge.

3. Water and Polarity Why do some chemicals dissolve in water?
Chapter 2
Water and Polarity
Why do some chemicals dissolve in water?
- The polar nature of water largely determines its solvent properties
 Ionic and poly substances are referred to “hydrophilic”
 Nonpolar molecules do not dissolve in water and are referred to “hydrophobic”

4. Water and Polarity Why do amphiphilic molecules form micelles?
Chapter 2
Water and Polarity
Why do amphiphilic molecules form micelles?
- A single molecule may have both polar (hydrophilic) and nonpolar (hydrophobic) portions  Amphiphilic

5. Chapter 2
Hydrogen Bonds
Why does water have such interesting and unique properties?
- There is an electrostatic attraction between the oxygen atom of one water molecule and the hydrogen of another, called a hydrogen bond.
 In liquid water, each molecule of water forms hydrogen bonds with an average of 3.4 other molecules.
 In ice, each molecule is fixed in space and forms hydrogen bonds with a full complement of four other molecules.

6. Hydrogen Bonds Hydrogen bonding gives unusual properties to water.
Chapter 2
Hydrogen Bonds
Hydrogen bonding gives unusual properties to water.
- Water has a higher melting point, boiling point, and heat of vaporization than most other common solvents because of its great internal cohesion.

7. Hydrogen Bonds Water forms hydrogen bonds with polar solutes
Chapter 2
Hydrogen Bonds
Water forms hydrogen bonds with polar solutes
- Hydrogen bonds readily form between an electronegative atom (the hydrogen acceptor, usually oxygen or nitrogen) and a hydrogen atom bound to another electronegative atom (the hydrogen donor)
- Hydrogen in C-H do not participate in hydrogen bonding

8. Hydrogen Bonds Water forms hydrogen bonds with polar solutes
Chapter 2
Hydrogen Bonds
Water forms hydrogen bonds with polar solutes
- Alcohols, aldehydes, ketones, and compounds containing N-H bonds all form hydrogen bonds with water molecules and tend to be soluble in water

9. Hydrogen Bonds Water interacts with electrostatically charged solutes
Chapter 2
Hydrogen Bonds
Water interacts with electrostatically charged solutes
- Water, a polar solvent, dissolves most biomolecules, which are generally charged or polar compounds: hydrophilic (Geek, “water-loving”)

10. Chapter 2
Hydrogen Bonds

11. HA  H+ + A-, Keq = [H+][A-]/[HA] = Ka, dissociation constant
Chapter 2
Acids, Base, and pH
What are acids and bases?
- Acids: Molecules that acts as a proton donor; Bases: a proton acceptor
 The stronger the acid, the greater its tendency to lose its proton
HA  H+ + A-, Keq = [H+][A-]/[HA] = Ka, dissociation constant
 Stronger acids have larger dissociation constants
pKa = log(1/ Ka) = -log Ka
 The stronger the tendency to dissociate a proton, the stronger is the acid and the lower its pKa

12. Equilibrium constant, Keq = [H+][OH-]/[H2O]
Chapter 2
Acids, Base, and pH
What is pH?
- Water molecules have a slight tendency to undergo reversible ionization
H2O  H+ + OH-
- The degree of ionization of water at equilibrium is small: ~ two of every 109 molecules
Equilibrium constant, Keq = [H+][OH-]/[H2O]
- [H+][OH-] = 1  M2
 Equal concentrations of H+ and OH-: neutral pH
 [H+][OH-] = [H+]2 = 1  M2  [H+] = 1  10-7 M
 As the ion product of water is constant, whenever [H]+ is greater than 1  10-7 M, [OH-] must become less than 1  10-7 M.

13. Acids, Base, and pH - pH = log(1/ [H+] ) = -log [H+]
Chapter 2
Acids, Base, and pH
- pH = log(1/ [H+] ) = -log [H+]
 The symbol p denotes “negative logarithm of” e.g., [H+] = 1  10-7 M  pH = 7
- The pH of an aqueous solution can be measured using
i) dyes such as litmus, phenolphthalein, and phenol red, which undergo color changes
as a proton dissociates from the dye molecule
ii) a electrode that is sensitive to [H+] but insensitive to Na+, K+, and other cations
- The pH affects the structure and activity of biological macromolecules, and it is often
used for medical diagnoses: e.g., diabetes (the pH of the blood plasma < 7.4)

14. Chapter 2
Acids, Base, and pH

15. Chapter 2
Acids, Base, and pH

16. Chapter 2
Acids, Base, and pH

17. Chapter 2
Titration Curves
- Titration is used to determine the amount of an acid in a given solution
 In general, the NaOH is added until the acid is consumed, as determined with an indicator dye or a pH meter
- A plot of pH against the amount of NaOH added reveals the pKa of the weak acids
 As NaOH is gradually introduced to the acid solution, the OH- combines with the free H+ to form H2O, satisfying the equilibrium relationship (H2O  H+ + OH-)
 As free H+ is removed, acids dissociate further to satisfy its own equilibrium
constant (dissociation constant) (HA  H+ + A-)
 At the end point, all the acid has lost its protons to OH-

18. Chapter 2
Titration Curves

19. Buffers How do buffers work?
Chapter 2
Buffers
How do buffers work?
A buffer solution consists of a mixture of a weak acid and its conjugate base.
Buffer solutions tend to resist a change in pH on the addition of moderate amounts of strong acid or base.
Henderson-Hasselbalch equation
pH = pKa + log[A-]/[HA]
- Useful in predicting the properties of buffer solutions used to control the pH of reaction mixtures.

20. Chapter 2
Buffers
- Almost every biological process is pH dependent; a small change in pH produces a large change in the rate of the process
 e.g., Enzymes and many of the molecules contain ionizable groups with characteristic pKa values
 Ionic interactions are among the forces that stabilize a protein molecule
- Cells and organisms maintain a specific and constant cytosolic pH
 Constancy of pH is achieved by biological buffers
- The intracellular and extracellular fluids of multicellular organisms have a characteristic and nearly constant pH, which is preserved by buffer systems
- Two especially important biological buffers: phosphate and bicarbonate systems

21. Chapter 2
Table 2-8, p. 60

22. Chapter 2
Summary

23. Chapter 2
Summary