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

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.

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”

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

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.

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.

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

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

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”)

Chapter 2 Hydrogen Bonds

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

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  10-14 M2  Equal concentrations of H+ and OH-: neutral pH  [H+][OH-] = [H+]2 = 1  10-14 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.

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)

Chapter 2 Acids, Base, and pH

Chapter 2 Acids, Base, and pH

Chapter 2 Acids, Base, and pH

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-

Chapter 2 Titration Curves

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.

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

Chapter 2 Table 2-8, p. 60

Chapter 2 Summary

Chapter 2 Summary