Water, acid base balance, buffer systems

Summarised lecture for undergraduates (beginners) in the course of Biochemistry

1. Water, pH scale Acid–Base Balance Buffers systems Dr. Male Keneth

2. Water 2

3. Water 65% of our bodies are water; cell are 70-95% water. 70% of our earth is ocean water. 3 different forms of water: Solid (ice), liquid (water), and gas (steam)

4. Water Electrons of each covalent bond are not shared equally between the H & O atoms. Oxygen has a stronger attraction. The unequal sharing causes the Oxygen to have a slightly negative (-) charge, while the hydrogen ends have a slightly positive (+) charge.

5. Polarity A molecule in which opposite ends have opposite electrical charges is called a POLAR molecule for example water.

6. Water Formula for water: H 2 O H atoms are joined to O atoms by a single covalent bond

7. Unusual and unique properties of water Boiling Point and Freezing Point Surface Tension , Heat of Vaporization , and Vapor Pressure Viscosity and Cohesion Solid State Liquid State Gas State Solvent

8. Boiling Point and Freezing Point Bp and Fp decrease as molecular size decreases, but not for water: Water requires more energy to break its hydrogen bonds before it boils/ freezes: Vital in living organisms living in water; as they could die instantly Also vital for cooling body temperature via sweating.

9. Surface Tension , Heat of Vaporization , and Vapor Pressure S.T: high, due to the hydrogen bonds. HoV : very high. V.P: inversely proportional to intermolecular forces: water has strong intermolecular forces = low V.P

10. Viscosity and Cohesion and Adhesion Viscosity : the property of fluid having high resistance to flow: water is viscous due to stronger intermolecular forces Cohesion: intermolecular forces between like molecules; this is why water molecules are able to hold themselves together in a drop. Due to its polarity. Importance : ????

11. Solid State and Liquid state Water is more dense than ice: why? As water cools below 4 C, hydrogen bonds rearrange themselves into an open crystalline, hexagonal structure, making water molecules to be held further apart Molecules are held tightly packed in water’s liquid state: being held by hydrogen bonds, moving freely. Why does this imply? Why is it important for living organisms?

12. Gas state As water boils, the hydrogen bonds are broken as particles move faster. Lack of hydrogen bonds explains why steam is causes worse burns water. All the energy to break the hydrogen bonds is contained in steam. And as stem is converted to liquid water, more heat energy is released.

13. Universal Solvent Because of it’s polarity, it dissociates many particles. Oxygen has a slightly negative charge, while the two hydrogens have a slightly positive charge. The slightly negative particles of a compound will be attracted to water's hydrogen atoms the slightly positive particles will be attracted to water's oxygen molecule; this causes the compound to dissociate

14. Water dissolves polar compounds 15 solvation shell or hydration shell

15. Hydrogen Bonding of Water 16 Crystal lattice of ice One H 2 O molecule can associate with 4 other H 2 0 molecules

16. Non-polar substances are insoluble in water 17 Many lipids are amphipathic

17. Micelles An aggregate of molecules in a colloidal solution, such as those formed by detergents. A typical micelle in aqueous solution forms an aggregate with the hydrophilic "head" regions in contact with surrounding solvent, sequestering the hydrophobic single-tail regions in the micelle center. 

18. How detergents work? 19

19. IONISATION OF WATER K eq , K w and pH H 2 O is the medium of biological systems. Dissociation of ions from biological molecules occurs in H 2 O. Water is essentially a neutral molecule but will ionize to a small degree. 20

20. Ionization of Water 21 K eq = [H + ] [OH - ] [H 2 O] H 2 0 H + + OH - K eq =1.8 X 10 -16 M [H 2 O] = 55.5 M [H 2 O] K eq = [H + ] [OH - ] (1.8 X 10 -16 M)(55.5 M ) = [H + ] [OH - ] 1.0 X 10 -14 M 2 = [H + ] [OH - ] = K w If [H + ]=[OH - ] then [H + ] = 1.0 X 10 -7 This equilibrium can be calculated as for any reaction:

21. Since the concentration of H 2 O is very high (55.5M) relative to [H + ] and [OH – ], it is generally removed from the equation by multiplying both sides by 55.5 yielding a new term, K w (ion product/ dissociation constant) : K w = [H + ][OH – ] In pure water, to which no acids or bases have been added: K w = 1 x 10 –14 M 2 at room temperature. Hence 10 –14 = [H + ][OH – ] 22

22. From 10 –14 = [H + ][OH – ] and taking log; –14 log 10 = log[H + ] + log[OH – ] -14 = log[H + ] + log[OH – ] 14 = -log[H + ] - log[OH – ] Therefore, 14 = pH + pOH At neutrality, [H + ] = [OH - ] = 1 x 10 –7 M Thus pH = 7, and also pOH = 7. Note: pH is the negative log of hydrogen ion concentration of any solution. 23 pH = –log[H + ] pOH = -log[OH – ]

23. pH Scale 24 Devised by Sorenson (1902) [H+] can range from 0M and 1 X 10 -14 M using a log scale simplifies notation pH = -log [H + ] Neutral pH = 7.0

24. Biological significance of pH H + is one of the most important ions in biological systems. Its conc. affects cellular and organismal processes as all metabolic processes are pH dependent. Structure and function of proteins depends on ionic interactions which are pH dependent. Movement of ions across membranes depends upon their net charge as determined by pH. The ionic state of the nucleic acids, lipids, mucopolysaccharides are determined by pH. All enzymes function best within optimum pH range. O 2 & CO 2 transport, release/gaseous exchange are pH dependent. The [H + ] also plays a role in energy generation and endocytosis. 25

25. Acid–Base Balance 28

26. Acid–Base Balance pH of body fluids is altered by the introduction of acids or bases Acids and bases may be strong or weak Strong acids dissociate completely (only HCl is relevant physiologically) Weak acids do not dissociate completely and thus affect the pH less compared to strong acids (e.g. carbonic acid) 29

27. Weak Acids and Bases Equilibria 30 Strong acids / bases – dissociate completely Weak acids / bases – dissociate only partially The pH depends on the degree of dissociation Enzyme activity is sensitive to pH Weak acid/bases play important role in protein structure/function.

28. K a In biological processes various weak acids and bases are encountered, e.g. the acidic and basic amino acids, nucleotides, phospholipids, etc. Weak acids and bases in solution do not fully dissociate and, therefore, there is an equilibrium between the acid and its conjugate base. This equilibrium can be calculated and is termed the equilibrium constant = K a . K a is also referred to as the dissociation constant . 31

29. Acid/conjugate base pairs 32 HA + H 2 O A - + H 3 O + HA A - + H + HA = acid ( donates H + ) A - = Conjugate base (accepts H + ) K a = [H + ][A - ] [HA] K a & pK a value describe tendency to loose H + large K a = stronger acid small K a = weaker acid pK a = - log K a

30. pKa values determined by titration 33

31. Phosphate has three ionizable H + and three pKas 34

32. Buffers Buffers are aqueous systems that resist changes in pH when small amounts of a strong acid or base are added. A buffered system consist of a weak acid and its conjugate base. Buffers are effective at pHs that are within +/-1 pH unit of the pKa The most effective buffering occurs at the region of minimum slope on a titration curve (i.e. around the pKa ). 35

33. Buffers At pK a the pH of a solution does not change appreciably even when large amounts of acid or base are added. This phenomenon is known as buffering . Dissolved compounds that stabilize pH by providing or removing H + Weak acids or weak bases that absorb or release H + are buffers 36

34. Buffer Systems Buffer System: consists of a combination of a weak acid and the anion released by its dissociation (its conjugate base) The anion functions as a weak base: H 2 CO 3 (acid)  H + + HCO 3 - (base) In solution, molecules of weak acid exist in equilibrium with its dissociation products (meaning all three species exist in plasma) 37

35. Henderson- Hasselbach Equation 38 1) K a = [H + ][A - ] [HA] 2) [H + ] = K a [HA] [A - ] 3) -log[H + ] = -log K a -log [HA] [A - ] 4) -log[H + ] = -log K a +log [A - ] [HA] HA = weak acid A - = Conjugate base * This equation describes the relationship between pH, pKa and buffer concentration 5) pH = pK a + log [A - ] / [HA] pH values of buffered solutions can be calculated using Henderson-Hasselbalch equation

36. Case where 10% acetate ion, 90% acetic acid 39 pH = pK a + log (90) ( 0.1 ) [0.9] pH = 4.76 + (-0.95) pH = 3.81

37. pH = pK a + log 10 [0.5 ] ¯¯¯¯¯¯¯¯¯¯ [0.5] pH = 4.76 + 0 pH = 4.76 = pK a Ie . at neutral point pH is equal to pKa Also, there is maximum buffering Case where 50% acetate ion 50% acetic acid 40

38. pH = pK a + log 10 [0.9 ] ¯¯¯¯¯¯¯¯¯¯ [0.1] pH = 4.76 + 0.95 pH = 5.71 Case where 90% acetate ion 10% acetic acid 41

39. Alkaline conditions pH = pK a + log 10 [0.99 ] ¯¯¯¯¯¯¯¯¯¯ [0.01] pH = 4.76 + 2.00 pH = 6.76 Acidic conditions pH = pK a + log 10 [0.01 ] ¯¯¯¯¯¯¯¯¯ [0.99] pH = 4.76 - 2.00 pH = 2.76 Cases when buffering fails 42

40. Question 1. Calculate the pH of a solution containing a mixture of 0.25M acetic acid and 0.1M sodium acetate. The pKa of acetic acid is 4.76. Question 2. Calculate the ratio of lactic acid to lactate required in a buffer system of pH 5. The pKa of lactic acid is 3.86. 43

41. Buffer Systems in Body Fluids 44 Figure 27–7

42. 3 Major Physiological Buffer Systems Protein buffer systems: Help regulate pH in extracellular fluid (ECF) and intracellular fluid (ICF) Interact extensively with other buffer systems Carbonic acid–bicarbonate buffer system: Most important buffer of blood (ECF) Phosphate buffer system: Buffers pH of ICF and urine 45

43. 1. Protein Buffer Systems Amino acids in protein buffer systems Depend on free and terminal amino acids Respond to pH changes by accepting or releasing H + If pH rises: carboxyl group of amino acid dissociates, acting as weak acid, releasing a hydrogen ion If pH drops: carboxylate ion and amino group act as weak bases accept H + form carboxyl group and amino ion 46

44. Protein buffer (most have pKa = 7.4) These include Hemoglobin, serum albumins and other plasma proteins , proteins in interstitial fluid and in the intracellular fluid (ICF) Several of these groups have pKa value around 7.4. Since proteins are present in significant concentrations in living organisms, they are important powerful buffers e.g. The Hemoglobin, most abundant molecule in red blood cells. Because of its structure and cellular concentration, Hemoglobin plays a major role in maintaining blood pH. 47

45. The Hemoglobin Buffer System CO 2 diffuses across RBC membrane: no transport mechanism required As carbonic acid dissociates: bicarbonate ions diffuse into plasma in exchange for chloride ions (chloride shift) Hydrogen ions are buffered by hemoglobin molecules The only intracellular buffer system with an immediate effect on ECF pH Helps prevent major changes in pH when plasma P CO 2 is rising or falling 48

46. 2. The Carbonic Acid–Bicarbonate Buffer System Formed by carbonic acid and its dissociation products . Prevents changes in pH caused by organic acids and fixed acids in ECF/blood H + generated by acid production combines with bicarbonate in the plasma. This forms carbonic acid, which dissociates into CO 2 which is breathed out 52

47. 53 It is a carbonic acid/bicarbonate (H 2 CO 3 /HCO 3 - ) buffer system. It is the most important buffer system in the body (plasma and ECF) despite the fact that bicarbonate has pKa of 6.1 far below the physiological pH 7.4 It is so important because it responds very fast to changes in plasma pH through loss of carbon dioxide in the lungs and bicarbonate in the urine through the kidneys. CO 2 reacts with water to form carbonic acid

48. From 7.4 = 6.1 + log [bicarbonate ] [Carbonic acid] The ratio of HCO 3 - to CO 2 required to maintain the blood pH of 7.4 is regulated: Carbon dioxide conc. is adjusted by changes in the rate of respiration Bicarbonate conc. is regulated by the kidneys; If the [HCO 3 - ] decreases, the kidneys remove H + from the blood triggering a shift to the right; increasing [HCO 3 - ]. When excess HCO 3 - ions are produced, they are excreted by the kidneys triggering a shift to the left. 54

49. Figure 27–9 The Carbonic Acid–Bicarbonate Buffer System 55

50. Limitations of the Carbonic Acid Buffer System Hard to protect ECF from changes in pH that result from elevated or depressed levels of CO 2 (because CO 2 is part of it) Functions only when respiratory system and respiratory control centers are working normally Ability to buffer acids is limited by availability of bicarbonate ions 56

51. The Phosphate Buffer System Consists of anion H 2 PO 4 — (a weak acid) Works like the carbonic acid–bicarbonate buffer system. Is important in buffering pH of ICF. 57

52. 3. Phosphate buffer (pKa = 7.2) Phosphate buffer consists of weak acid conjugate base pair (H 2 PO 4 - / HPO 4 2- ) H 2 PO 4 -  H + + HPO 4 2- Dihydrogen phosphate Hydrogen phosphate It has a pKa of 7.2, the blood pH is 7.4, close to the phosphate buffer pKa of 7.2: looks a perfect buffer but the concentration of H 2 PO 4 - and HPO 4 2- in blood are too low to have a major effect. However, the concentration of phosphate buffer in intracellular fluid is higher (approx. 75mEq/L) than it is in blood (4mEq/L) It is therefore an important buffer in intracellular fluid (ICF). 58

53. Problems with Buffer Systems Provide only temporary solution to acid–base imbalance. Do not eliminate H + ions. Supply of buffer molecules is limited. 59

54. Maintenance of Acid–Base Balance Requires balancing H + gains and losses For homeostasis to be preserved, captured H + must either be: permanently tied up in water molecules through CO 2 removal at lungs OR removed from body fluids through secretion at kidney Thus, problems with either of these organs cause problems with acid/base balance Coordinates actions of buffer systems with: respiratory mechanisms renal mechanisms 60