1. CH339K Bonding, Water, Acids, Bases, and Buffers

2. Bonding Covalent Ionic Dipole Interactions Van der Waals Forces Hydrogen Bonds Hydrophobic Interactions

3. Covalent Bonds Electrons form new orbitals around multiple atomic nuclei Bond energy results from electrostatic force between redefined electron cloud and nuclei Strong – typically 150 – 450 kJ/mol

4. Common Covalent Bond Numbers For Biochemically Significant Elements AtomBond Number* C4 H1 O2 N3 P3,5 S2,4,6 *the bond numbers illustrated are typical of biological systems and should not be considered set in stone. i.e. Don’t depend on this for your Inorganic class!

5. Ionic Interactions Ribbon representation (A) and surface charge distribution of (B) of the lowest-energy solution structure of the SARS-CoV X4 ectodomain. Red and blue surfaces represent negative and positive electrostatic potentials, respectively. Biomolecules frequently have large numbers of charged groups Charge-charge interactions stabilize intra- and intermolecular structures

6. Ionic Interactions Energy from non-directional electrostatic force between ions Coulomb’s Law: Energy drops off as function of distance between charges (i.e. operates over long ranges) –  is the dielectric constant of the medium – k is the Permittivity of the Vacuum – sort of an absolute dielectric constant to which other dielectric constants relate. Ionic interactions tend to be quite strong.

7. Some enchanted evening, you may see a stranger; you may see a stranger across a … WEAK Dielectric STRONG Dielectric …empty room …crowded room

8. Common Dielectric Constants Namedielectric constant water80 methanol33 ethanol24.3 1-propanol20.1 1-butanol17.8 formic acid58 acetic acid6.15 acetone20.7 hexane2.02 benzene2.28

9. Dipoles Fixed dipoles –Molecules with asymmetric charge distributions form dipoles Induced Dipoles –One dipole can induce a charge in an adjacent molecule

10. Dipole moments  = q·x –Where  is the dipole moment –q is the charge –x is the distance between the charges The larger the dipole moment, the more polar the molecule.

12. van der Waals Interactions Technically, all induced dipole interactions are van der Waals interactions Biochemists usually mean induced dipole-induced dipole (London Dispersion) forces Any atom will have an uneven distribution of charge at any given instant

13. Van der Waals (cont.) That temporary dipole will induce a dipole in adjacent atoms This results in a net attractive force between atoms Force is weak -.5 to 2 kJ/mol Net biochemical effect – molecules that FIT together STICK together.

14. Van der Waals (cont.) If you live in Central Texas, you see van der Waals forces in action every summer night:

15. Artificial Geckos Climb building Crawl through window Find target Detonate Real gecko toe hair (Courtesy of Geico) Synthetic gecko toe hair

16. Protein – DNA Interaction

17. Charge Interaction Energies and Distance

18. Van der Waals Forces (cont.) London Dispersion Forces cause particles to come together… Until they get too close.

19. Atomic Radii The Lennard-Jones potential describes the interaction between a pair of neutral atoms: London Dispersion ForceElectron Cloud Repulsion

20. Multiple molecular contacts can mediate binding hexokinase Energy of van der Waals contacts can subsidize conformational changes in molecules.

21. Hydrogen Bonds Hydrogen Bonds form between –A hydrogen covalently bound to an electronegative atom –Another electronegative atom

22. Hydrogen Bonds (cont.) The group to which the hydrogen is covalently bound is the donor. The other group is the acceptor. Donors: –-OH, -NH 2, -SH (lesser donor) Acceptors –-N:, =O:, -O: H

23. Hydrogen Bonds (cont.) Bond length is less than vdW contact distances distance between the nuclei of the hydrogen bond acceptor and the hydrogen itself can be as short as 1.8-1.9 Å, well below the sum of the atomic radii (e.g. 1.2Å for hydrogen and ~1.5Å for oxygen and nitrogen)

24. Hydrogen Bonds (cont.) Intermediate strength: 5 – 30 kJ/mol Hydrogen bonds are not just electrostatic – partially covalent Therefore, they are directional

25.  -helix: An internal protein structure mediated by Hydrogen Bonds (amide hydrogens to carbonyl oxygens)

26. Binding from both H-Bonds and vdW Contacts EcoR1 – a DNA-cleaving Protein DNA Double Helix

27. Recap of Bond Energies (typical) STRENGTH (kcal/mole) BOND TYPELENGTH (nm)IN VACUUMIN WATER Covalent0.1590 Noncovalent: ionic0.25803 hydrogen0.3041 van der Waals attraction (per atom)0.350.1 Copyright © 2002 Bruce Alberts, Dennis Bray, Julian Lewis, Martin Raff, Keith Roberts, and James D. WatsonBruce Alberts, Dennis Bray, Julian Lewis, Martin Raff, Keith Roberts, and James D. Watson

28. Water Structure

29. Water Forms Clusters in Solution

30. Dissolving nonpolar molecules Solvating a nonpolar molecule imposes order on the surrounding water -  S < 0) Clathrate cage of ordered water

31. Hydrophobic interactions Solvating a non-polar material in water decreases entropy – forces water into an ordered structure Minimal energy is when water is least ordered – the more you can pack non- polar materials, the less surface area exposed to solvent.

32. Hydrophobic Effect

33. Recap – bonding in biomolecules Aka Salt Links

34. Ice Floats

35. Water –Has a high specific heat –Has a high heat of vaporization –Is an excellent solvent for polar materials –Is a powerful dielectric –Readily forms hydrogen bonds –Has a strong surface tension –Is less dense when it freezes (i.e. ice floats)

36. Acids and Bases Definitions –Arrhenius Acids are substances which produce an excess of H + ions in water (HCl) Bases are substances which produce an excess of OH - in water (NaOH) –Bronsted-Lowry Acids are substances which can donate a proton in a chemical reaction. (HF) Bases are substances which can accept a proton in a chemical reaction (NH 3 ) –Lewis Acids are electron - pair acceptors.(BF 3 ) Bases are electron - pair donors (CaO)

37. Conjugate Pairs Every acid has its conjugate base Every base has its conjugate acid Conjugate AcidConjugate Base H 3 C - COOHH 3 C-COO - NH 4 + NH 3 Water can act as both an acid and a base.

38. Ionization of Water

39. The extra proton of a hydronium ion is not confined to a single water molecule.

40. Ionization of Water K w = [H + ] [OH - ] = 10 -14 K w is small, so water has a very low amount of ionization. (at 25 C, 1 atm) K w is temperature dependent! K w = 10 -13 at 60 C. [H + ] is a measure of acidity. However, it ranges over many orders of magnitude. For ease of arithmetic, we generally refer to pH pH = -log[H + ] (for our purposes)

41. Typical pH Values SubstancepH Stomach acid1.5 - 2.5 Coca-cola2.5 Human saliva6.5 Human blood7.5 Human urine5 - 8 Oven cleaner14

42. Acids and Bases A Strong acid is one that dissociates completely in water; a weak acid is one that doesn’t. –Hydrochloric, Hydroiodic, Hydrobromic, Nitric, Sulfuric, Perchloric All biochemically significant acids and bases are weak – i.e. they don’t dissociate completely (except for HCl – stomach acid)

43. Acids and Bases The strength of a weak acid can be described by the Ka (ion product) and the pKa. The lower the pKa, the stronger the acid.

44. Typical Ka’s and pKa’s AcidKapKa Acetic1.8 x 10 -5 4.76 Formic1.7 x 10 -4 3.75 Benzoic6.5 x 10 -5 4.19 Carbonic4.3 x 10 -7 6.37 Imidazole2.8 x 10 -7 6.55 Phenol1.3 x 10 -10 9.89

45. pH for Strong Acids Since a strong acid dissociates completely: pH = -log([Acid]) For a 0.1 M (100 mM) solution of HCl: pH = -log(0.1) = 1 Well, that was difficult…

46. pH for Weak Acids Depends on the Ka What’s the pH of a 100 mM solution of Acetic Acid? [H+] = 0.00134 M

47. Shortcut The quadratic solution is a pain, but we can approximate: Accurate as long as acid < 5% dissociated [H+] = 0.00134 M

48. Titrating a Strong Acid 10 ml of an HCl sln. Titrate with 0.5 M NaOH OH - + H + → H 2 O Takes 8.5 ml NaOH to bring solution to neutrality

49. Titrating a Weak Acid Titrating.1 M HAc Initial pH is 2.88 instead of 1 Little change until large amounts of NaOH have been added Buffering effect Caused by equilibrium that exists between a weak acid and conjugate base.

50. Henderson-Hasselbalch Equation

51. Predicting pH Let’s make 1 liter of a solution that is 0.1 M in acetic acid ( pKa = 4.76 ) and 0.3 M in sodium acetate.

52. Buffering Effect pH depends on pKa and the ratio of conjugate base to acid. Addition of significant amounts of acid or base changes the ratio of conjugate base to conjugate acid pH changes as the log of that ratio Result is resistance to pH change in a buffered solution

53. Factors impacting pKa: Ionic Strength The ionic strength of a system is the sum of contributions from all ions present: where C i is the concentration of ion I, Z i is the charge on ion I

54. Factors impacting pKa: Ionic Strength Example: Phosphoric Acid has 3 pKa’s H 3 PO 4 ⇄ H + + H 2 PO 4 - ⇄ 2H + + HPO 4 -2 ⇄ 3H + + PO 4 -3 pKa 1 pKa 2 pKa 3 pKa 2 = 7.2 at ionic strength J = 0 pKa 2 = 6.86 at physiological ionic strengths (Physiological saline is 0.91% NaCl. Calculation of J is left as an exercise for the student)

55. Factors impacting pKa: Temperature pKa can (i.e. does) vary with temperature Example: one of the most common biochemical buffers is Tris (tris(hydroxymethyl)aminomethane) Tris is a good buffer at near- physiological pHs, is biologically pretty inert, and is (relatively) inexpensive. BUT Tris has a large thermal coefficient: - 0.031 units/ o C At 25 o C,pKa = 8.30 At 0 o CpKa = 7.77

56. A Physiological Example: Blood pH Blood pH is maintained at ~7.4 –pH below 7.35 is acidosis –pH above 7.45 is alkalosis pH 7.8 is generally fatal

57. Blood pH Control Blood pH is regulated by four buffer systems: CarbonateH 2 CO 3 ⇄ H + + HCO 3 - pKa = 6.1 PhosphateH 2 PO 4 - ⇄ H + + HPO 4 -2 pKa = 6.9 Plasma Proteins Hemoglobin The primary system, carbonate, has 3 interlocking equilibria: CO 2 (g) ⇄ CO 2 (aq) + H 2 O ⇄ H 2 CO 3 ⇄ H + + HCO 3 - Excess H + or HCO 3 - drives the equilibrium to the left Excess H 2 CO 3 drives the equilibrium to the right

58. Blood pH Control CO 2 (g) ⇄ CO 2 (aq) + H 2 O ⇄ H 2 CO 3 ⇄ H + + HCO 3 - Diseases that effect the level of [HCO 3 - ] are metabolic effects, due to changes in cellular metabolism. Diseases that change [H 2 CO 3 ] are respiratory effects; the lungs control the exchange of CO 2, and therefore the concentration of H 2 CO 3.

59. Blood pH Control Metabolic Acidosis: Diseases such as diabetes or diarrhea result in an excess of H + in the tissues. [HCO 3 - ] goes DOWN (equilibrium pushed to left) Blood pH goes DOWN. (equilibrium to left; higher carbonic acid, lower bicarbonate) CO 2 (g) ⇄ CO 2 (aq) + H 2 O ⇄ H 2 CO 3 ⇄ H + + HCO 3 -

60. Blood pH Control Metabolic Alkalosis: Vomiting (intoxication, gastrointestinal illnesses) causes loss of H+. [HCO 3 - ] goes UP (equilibrium pulled to right) Blood pH goes UP. (equilibrium to right; lowerer carbonic acid, higher bicarbonate) CO 2 (g) ⇄ CO 2 (aq) + H 2 O ⇄ H 2 CO 3 ⇄ H + + HCO 3 -

61. Blood pH Control Respiratory Acidosis: In conditions like emphysema, pneumonia, your lungs do not work effectively to clear CO 2. [H 2 CO 3 ] goes UP (driven by carbon dioxide build-up.) Blood pH goes DOWN (as carbonic acid accumulates.) CO 2 (g) ⇄ CO 2 (aq) + H 2 O ⇄ H 2 CO 3 ⇄ H + + HCO 3 -

62. Blood pH Control Respiratory Alkalosis: When you hyperventilate or become hysterical, you blow off lots of CO 2. [H 2 CO 3 ] goes DOWN (since its being withdrawn as CO 2.) Blood pH goes UP (less carbonic acid.) CO 2 (g) ⇄ CO 2 (aq) + H 2 O ⇄ H 2 CO 3 ⇄ H + + HCO 3 -

63. A Practical Buffer Problem Benzoic acid is a weak carboxylic acid that is reasonably soluble in water (3.4 g/l). –Molecular Weight:122.12 g/mol –pK a 4.21 I wish to make 4 liters of 10 mM Sodium Benzoate buffer, pH 5.0. I have solid benzoic acid in a jar, a stock solution of 5 M Sodium Hydroxide (NaOH), a 4 liter graduated cylinder and all the deionized, distilled water I can use. How do I make the buffer?

64. Practical Buffer Problem (cont.) Step 1: Okay, how much benzoic acid do I need? Since benzoate will be the buffering ion, I want my solution to be 10 mM in total benzoate. Solid benzoic acid is my only source of benzoate, so I need to add 10 mM worth:

65. Practical Buffer Problem (cont.) Step 2: How do I get it to the right pH? The conjugate base of benzoic acid is benzoate anion. Addition of a strong base (like NaOH) to benzoic acid converts it to benzoate. The pH of the solution depends on the ratio of conjugate base to conjugate acid as determined by the Henderson-Hasselbach equation. How much benzoic acid to I have to convert to benzoate base to give me the desired ratio of conjugate base to conjugate acid?

66. Practical Buffer Problem (cont.) Rats! 1 equation with 2 unknowns…

67. But wait! That’s not all! We also know that total benzoate is 10 mM We need to convert 8.61 mM benzoic acid to the conjugate base, benzoate. To convert 8.61 mM benzoic acid to 8.61 mM benzoate, we need to add 8.61 mM (.00861 M) NaOH So: add 4.88 g of Benzoic Acid, 6.89 ml of 5M NaOH, and enough H 2 O to make 4 liters.

68. What you ought to be able to do after this agony Know the different “bond” types involved in biomolecular interactions Know their relative strengths (i.e. bond energies) Be able to calculate ionic interaction energies Understand what the van der Waals radius is Be able to recognize H-bond acceptors and donors Understand hydrophobic interactions Be able to identify conjugate acids and bases Understand the ideas of pH, Ka, and pKa Be able to calculate pH and to calculate how to make buffers