Chapter 2 Aqueous Chemistry

1. Essential Biochemistry Third Edition Charlotte W. Pratt | Kathleen Cornely Chapter 2 Aqueous Chemistry

2. KEY CONCEPTS: Section 2-1 • The polar water molecule forms hydrogen bonds with other molecules. • Noncovalent forces, including H bonds, ionic interactions, and van der Waals forces, act on biological molecules. • Water dissolves both ionic and polar substances.

3. Water is abundant in living organisms. About 60% of humans is water.

4. Water is an effective polar solvent. • Water has tetrahedral geometry because of its electronic structure. • Water is polar. An electron from hydrogen is shared with one from oxygen. Oxygen has 2 lone pair electrons.

5. Opposite charges attract. • Water molecules orient so that oxygens are positioned near hydrogen(s) in a neighboring molecule. • Oxygen has a partial negative charge; hydrogen has a partial positive charge.

6. Hydrogen bonds are prevalent among biological molecules. • A hydrogen bond is a noncovalent interaction between a molecule containing an electronegative atom (such as O or N) and another molecule containing an electropositive hydrogen.

7. Hydrogen bonding capability is based on electronegativity. Electronegativity is a periodic function. The farther the element is to the right of the Periodic Table, the more electronegative it is.

8. Hydrogen bonds can form between molecules other than two waters. Important Observation:The hydroxyl functional group and the amine groups are key players in hydrogen bonding in biomolecules.

9. Example: Hydrogen Bonds in DNA Oxygen Nitrogen Hydrogen Cytosine and Guanine are bases in DNA that form hydrogen bonds. Notice how hydrogen is always bonding with an electronegative atom.

10. Atomic distance matters. • When the atoms are too close, their orbitals overlap. • When the orbitals are as close as possible without overlapping, a hydrogen bond can form. • Hydrogen bond distance is 2.7Å on average.

11. Forces are involved in biomolecular structures • Sometimes relatively weak forces play an important role in biomolecular structures. • Why? Because there are many of them!

12. Analogy: Gulliver’s Travels Gulliver was restrained by the minute Lilliputians just as biomolecules are restrained by weak forces.

13. KEY CONCEPTS: Section 2-2 • The hydrophobic effect, which is driven by entropy, excludes nonpolar substances from water. • Amphiphilic molecules form micelles or bilayers.

14. What happens when oil and water mix? • Nonpolar molecules such as oil are not miscible with polar solvents such as water.

15. Nonpolar molecules tend to aggregate in polar solvents UnfavorableMany H2O molecules are ordered around the nonpolar molecules PreferredFewer H2O molecules are ordered around the nonpolar molecules Low Entropy(lower disorder of water) High Entropy(higher disorder of water) The hydrophobic effect is the phenomenon by which nonpolar molecules aggregate to avoid contact with hydrophilic molecules, particularly water.

16. Most lipids are amphiphilic. • Fatty acids are a type of lipid composed of a hydrophilic head group and a hydrophobic (nonpolar) hydrocarbon tail – the definition of an amphiphile.

17. Amphiphilic molecules form micelles in aqueous solutions. In H2O View in 2D

18. Amphiphilic molecules form vesicles in aqueous solutions. Forming a lipid bilayer vesicle in aqueous solution Cross-sectional view in 3D

19. KEY CONCEPTS: Section 2-3 • Water ionizes to form H+ and OH-. • An acid’s pK value describes its tendency to ionize. • The pH of a solution of acid depends on the pK and the concentrations of the acid and its conjugate base.

20. Recall From General Chemistry:Water Ionizes OH– 2 H2O H3O+ + Hydroxide ion Hydronium ion More simply: H2O H+ + OH– [H+] x [OH–] K = [H2O]

21. Equilibrium Constant for Ionization of Water [H+] x [OH–] K = [H2O] Becomes 1 H2O H+ + OH– Here, K = Kw = ionization constant for water, which has a known value. [H+] x [OH–] = 10–14 Kw =

22. Defining pH [H+] x [OH–] = 10–14 Kw = • In pure water, [H+] = [OH–] = 10–7 M • Pure water is neutral. • Solutions in which [H+] > 10–7M = acidic • Solutions in which [H+] < 10–7M = basic • An easier approach: pH = –log [H+]

23. The pH Scale • As [H+] increases, [OH–] decreases. • pH is logarithmically proportional to [H+].

24. pK describes an acid’s tendency to ionize. • Example: acetic acid CH3COOH CH3COO– + H+ [CH3COO–] x [H+] K = = 1.74 x 10–5 [CH3COOH] pK = –log (1.74 x 10–5) = 4.76

25. Polyprotic acids ionize more than once and have multiple pKs. • H3PO4 = phosphoric acid pK = 2.15 • H2PO4- = “monobasic” pK = 6.82 (or 7.21) • HPO42- = “dibasic” pK = 12.38 • PO43- = “tribasic” + H+ + H+ + H+ H3PO4 H2PO4-HPO42-PO43-

26. pH is related to pK. HA H+ + A– Take negative log on both sides. [H+] x [A–] K = – log K = – log [H+] – log [HA] pK = pH – log [A–] [A–] [A–] Solving for pH gives theHenderson-Hasselbalch Equation. [HA] [HA] [HA] pH = pK + log

27. Henderson-Hasselbalch Equation pH = pK + log OR Change the sign Invert the ratio [HA] pH = pK – log [A–] [A–] [HA]

28. pK and pH convey protonation state. • Example: a generic amino acid At pH’s < 3.5, amino acids have both protons and a net positive charge – COOH and NH3+. pK = 3.5 When 3.5 < pH < 9.0, amino acidslose the proton on the carboxyl carbon and are zwitterions; charge is neutral. At pH’s > 9.0, amino acidshave lost both protons and have a net negative charge – COO– and NH2. pK = 9.0

29. Tools and Techniques • Buffer solutions resist changes in pH.

30. Convenient Rule of ThumbpH = pK ± 1 • When [A–] = 10 [HA] • When [HA] = 10 [A-] pH = pK + log (10) = pK+ 1 pH = pK + log (0.1) = pK– 1 Buffers tend to resist changes in pH when [HA] and [A–] differ by no more than a factor of 10.