Major concepts

1. What do ΔG° and K_eq tell us about a reaction?
2. How do water molecules relate to each other and to other molecules? How do these relationships affect cells and biochemical interactions within cells?
3. What happens to a weak acid in solution and to the pH of the solution when base is added?

Core knowledge

1. What is the equation that relates ΔG, ΔG°, and K_eq
2. What are the different weak interactions that are important in biochemistry?
3. What is pH? What is the pK_a of a weak acid? What is the Henderson-Hasselbalch equation?

Molecules in biochemistry
A. Review functional groups (organic) if necessary - Fig. 1-15
B. Macromolecules are composed of smaller molecules (subunits)
     1. Types = proteins, nucleic acids, polysaccharides, and some lipids
     2. Subunits = amino acids, nucleotides, sugars, and, mostly, fatty acids
     3. Three-dimensional structure: configuration and conformation

Quick review of free energy and equilibrium (thermodynamics)
A. Living systems are not at equilibrium; that is, things always change.
      This involves energy. - Fig. 1-24
B. ΔG = ΔH - T ΔS - review the terms in this equation if necessary
C. Because ΔH in cells is usually small, biochemists talk about exergonic and
     endergonic reactions, not exothermic and endothermic reactions. - Fig. 1-26
D. ΔG for coupled reactions is additive.
E. ΔG for a reaction is related to K_eqfor the reaction.
     For a reaction in which A + B ↔ C + D, K_eq= ([C] [D]) / ([A] [B])
     ΔG° = ΔG measured at standard temperature (273 K) and pressure (1 atm)
     ΔG = ΔG°′ + R T ln ([C] [D]) / ([A] [B]) = ΔG°′ + R T ln K_eq
F. At equilibrium, ΔG = 0, and ΔG°′ = - R T ln K_eq.

Water and weak interactions with biochemical molecules
A. Water molecules are polar and form hydrogen bonds with each other - Fig. 2-1
     Hydrogen bond: a specific attraction, represented by - - - ,
     between H with a covalent bond to O or N (H donor) and another O or N (H acceptor)
    This affects T_M, T_B, and heat of vaporization - Table 2-1.
B. Water forms hydrogen bonds with polar molecules - Fig. 2-3 and Fig. 2-5.
C. Water has electrostatic (ionic) interactions with charged solutes - Fig. 2-6
     Ionic interaction: interaction between two charged groups (ions, dipoles, and/or
                         groups with partial charges)
        Represented by positioning the ion and group near each other physically (no - - - )
D. Note these differences:
     1. Between hydrogen bonds and electrostatic interactions
        Hydrogen bonds are directional (more limited), longer, weaker.
     2. Between ionic bonds and electrostatic interactions
        Ionic bonds are between ions in a lattice, shorter, stronger
E. Hydration of polar and charged solutes is favorable; ΔS is positive.
F. Water does not form hydrogen bonds or electrostatic interactions with nonpolar solutes.
     In their presence, water molecules become more ordered - Fig. 2-7
G. Hydrophobic interactions occur between nonpolar molecules in the presence of water.
     Fig. 2-7

Van der Waals interactions
A. Definition: temporary attraction between the electron(s) of 1 atom and the nucleus of another
B. Strength and specificity: weakest, longest, least specific

Weak interactions in biochemistry
A. Be able to characterize different types of molecules - Table 2-2
B. Know the types and specificity of weak interactions for different types of molecules -  
      Table 2-6
C. Weak interactions are additive.

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Water and pH
A. Ionization of water - Fig. 2-14
     1. Reaction: 2 H₂O ↔ H₃O⁺ + OH^-    or H₂O ↔ H⁺ + OH^-
     2. K_eq = ([H⁺][OH^-])/[H₂O] = ([H⁺][OH^-])/55.5 M
     3. K_w = 55.5 M K_eq = [H⁺][OH^-] = 1 x 10^-14
B. pH = negative log of [H⁺]
C. neutral pH (pH 7) has [H⁺] = [OH^-] = 1 x 10^-7 M
D. Acid solutions have pH < 7, while basic solutions have pH > 7. Fig. 2-15

Strong and weak acids
A. Acids donate H⁺ and increase [H⁺]: HA (aq) ↔ H⁺ (aq) + A^- (aq)
      HA = the acid, and A^- = the conjugate base
B. Strong acids dissociate completely in solution: HCl, HNO₃, H₂SO₄.
C. Weak acids dissociate, but only when [H⁺] is at a certain level.
D. K_eq for an acid = ([H⁺] [A^-]) / [HA] = K_a. K_a is large for strong acids.
E. The negative log of K_a = pK_a(= log 1/K_a)
      pK_ais low for strong acids and higher for weak acids.
F. The pK_afor an acid is the pH at which [HA] = [A^-].
     In a titration of a weak acid, the pK_ais at the half-way point of the titration. Fig. 2-17
     At that point, the addition of base doesn't cause a significant change in pH, because
     there is a significant concentration of HA that can donate H⁺.

Strong and weak bases
A. Bases accept H⁺ and decrease [H⁺] in solution: B^- (aq) + H⁺ (aq) ↔ BH (aq)
B. Biochemists simplify discussion of bases by treating BH as a weak acid, so
     a weak base has a K_a and a pK_a just as a weak acid does.

Polyprotic acids donate more than one H⁺ and have complex titration curves. Fig. 2-16

Buffers = weak acids and their conjugate bases (HA + A^-)
A. When H⁺ is added to a buffer, it reacts with A^- → HA, and total [H⁺] remains constant.
       Therefore the pH remains the same.
B. When OH^- is added to a buffer, it reacts with HA → H₂O + A^-, and total [H⁺] remains
     constant.
C. This is true for the pK_a + 1 pH unit. Fig. 2-17, Fig. 2-18
D. The Henderson-Hasselbalch equation describes the behavior of buffers:
     pH = pK_a+ log ([A^-]/[HA])

Important examples of buffers in biochemistry
A. Carbonic acid and bicarbonate
     1. Carbonic acid is formed from the reaction of CO₂ + H₂O ↔ H₂CO₃
     2. The pK_aof carbonic acid is 3.77
B. Dihydrogen phosphate and hydrogen phosphate
     1. H₂PO₄^- ↔ HPO₄^2- + H⁺
     2. The pK_afor dihydrogen phosphate is 6.86.
C. Other common acids:
      1. Carboxylic acids: R-COOH ↔ R-COO^- + H⁺
         pK_a's of these acids are determined by R.
     2. Amines: R-NH₃⁺ ↔ R-NH₂ + H⁺
         pK_a's are also affected by R.

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