Water

1. Water

2. Water • Water is the most common chemical component of living organisms. • A 70 kg man contain approx 45 liters of water, maintenances of the amount depends on balance between intake & loss .

3. Normal daily water loss in sweat & expired air amounts to about 900 ml. • Over half of the body water is inside cells. The extracellular water is about 15 – 20 percent in the plasma. The remainder makes up the extra vascular, extracellular interstitial fluid.

4. Water has the ability to solvate organic & inorganic molecules. • It is due to its dipolar structure and its capability of forming hydrogen bond.

5. Water has two lone pairs of electorns bear a partial negative charge, is an excellent nucleophile . It acts as reactant and also as a product. • Nucleophilic attack by water results in the cleavage of amide, glycosides or ester bonds.

6. H₂O OH¯ + H⁺ • 2H₂O H₃O¯ + H⁺ • Water in the body is regulated by hypothalamic mechanisms that control thirst. Ant diuretic hormone (ADH) is involved in the retention or excretion of water by kidneys. • Water molecule is an irregular, slightly skewed tetrahedron with oxygen at its center.

7. O • 105° • H H • It has 105 degree angle between the hydrogen atoms. • Water is a dipole molecule with electrical charge distributed asymmetrically about its structure. • O¯δ • ⁺ δ H H ⁺ δ

8. The strongly electronegative oxygen pulls electrons away from the hydrogen nuclei leaving them with a partial positive charge. • Its two unshared electron pairs produce a local negative charge on oxygen. • Water molecules have a high dipole moment because of the polarity of the bonds.

9. Neighboring liquid water molecules interact with one another. • The intermolecular bonding between water molecules arises from the attraction between the partial negative charge on oxygen atom and the partial positive charge on the hydrogen atom of adjacent water molecules. This type of interaction is called hydrogen bonding.

10. H H • O¯δ • H⁺ δ • O¯δ • H ⁺ δ

11. Hydrogen bonding favors the self association • of water molecules. • Hydrogen bonding accounts for its high • viscosity, surface tension & boiling point. • Each molecule in liquid water associates • through hydrogen bond.

12. These bonds are relatively weak and transient • Rupture of a hydrogen bond in liquid water require, 4.5 k cal/mol. • Hydrogen Bonding enables water to dissolve many organic molecules which can participate in hydrogen bonding.

13. The oxygen atoms of aldehydes, ketones and amides provide lone pairs of electrons that can serve as hydrogen acceptor. • Alcohols and amines can serve both as hydrogen acceptors & as a hydrogen donors & thus they form hydrogen bond.

14. In proteins, Nucleic acids, and water, hydrogen bonds are essential to stabilize overall structure. • Covalent bond are the strongest force that hold molecules together. • Non covalent Forces make significant contributions to the structure. Stability and functions of macromolecules in living cells

15. Physical Properties Of Water • Physical properties of water suits to biological systems, include melting point boiling point. • Heat of vaporization (quantity of heat energy required to convert 1 gm of liquid to vapour).

16. Heat of fusion (quantity to heat energy required to transform 1 gm of solid to liquid at the melting point). • Specific heat (Amount of heat required to raise the temperature of 1 gm of substance by 1°C).

17. All these values of water are higher from other lower molecular weight substances due to strong intermolecular hydrogen bonding. • This property maintains the optimal body temp.

18. Van der Waals Forces • Van derwaals forces arises from attractions between transient dipole, generated by the rapid movement of electrons.

19. Multiple Forces Stabilize Biomolecules • DNA double helix structure is stabilized by hydrogen bonds • Each individual DNA strand is held together by covalent bond. • The two strands are held together by noncovalent interactions.

20. These non covalent interactions are, hydrogen bonds between nucleotide bases. • Vander Waals interactions between purine & Pyrimidine bases.

21. Dissociation of water • Water molecules exhibit a slight tendency of dissociation. • So water can act both as an acid and a base. • H2O + H2O H3O+ + OH- • A proton in aqueous solution in very mobile, going from water molecule to another with a period of about 10-15 seconds.

22. For practical purposes, the equation is written in a simple form. • H2O H+ + OH- • The proton do not exists in aqueous solution as a free ion but it is always associated with water molecule. So here we are really referring a hydrated proton. • [H+] [OH-] = Kw = 1 x 10-14 M2 • M= moles / L

23. For pure water • [H+] = [OH] = 1 x 10-7 M. • And the solution is said to be neutral neither acidic nor basic • The ion products depends on temperature • For example in human body at 370C the conc: of H+ & OH- in neutral solution will be 1.6 x 10 - 7 M.

24. pH • pH is the negative log of the hydrogen ion concentration. • pH = - log [H+] • The pure water at 25 0c • pH = -log [H+]= - log 10-7 = (-7)= 7.0 • Kw = [H+] [OH-] = 10-14 • log [H+] + log [OH-] = log 10-14 • pH+pOH = 14

25. Low pH values correspond to high concentration of H+. • High pH values correspond to low concentration of H+. • Acids are proton donors and bases are proton acceptors.

26. Strong acids. (Hcl, H2So4) and strong base (koH & NaoH) compeltely dissociate into aninos & cations in acidic solution. • Weak acids and weak bases, dissociates only partially in acidic solution .

27. Organic Compounds in living organisms exist in charged state. • Positively charged cations • Negatively charged Anions The presence of ionic forms in determined by hydrogen ion concentration [H⁺] expressed in terms of pH.

28. Because of the control of pH of organic ions living organisms are capable of preventing excessive charges in the intra and extracellular fluids & pH is usually maintained. The physiological pH is approx 7. pH is the negative log of [H⁺]

29. H₂ODissociationOH¯ + H⁺ H⁺ bond to the oxygen atom of an undissociated H₂O molecule to form hydronium ion [H₃O⁺]. H₂O + H₂O H₃O¯ + H⁺ So H₂O acts as an acid (donor of H⁺ donor or proton donor). So in this way • Acids are electron pair acceptor or H⁺ donor. • Bases are electron pair donors or H⁺ acceptors.

30. The equilibrium Constant K = [H⁺][OH¯] [ ] = molar conc. [H₂O] of anions K can be determined by measurement of electrical conductivity of pure water, which has the value K= 1.8 × 10⁻16 M (Molar)= moles/L Which is a very small ion conc. H₂O = 2+16

31. Molecular weight of H₂O = 18 Since 1 L H₂O weight is 1000 gm and 1 mol of water’s weight is 18 gm.

32. Molar conc. = Weight molecular weight The molar concentration of pure water is 55.5 M. Substitution of K & [H₂O] in the equation K = [H⁺][OH¯] [H₂O] [H⁺][OH¯] = [55.5] × (1.8 × 10⁻16M) [H⁺][OH¯] = 1.0 × 10⁻14(M/L)² = Kw. Kw is known as the ion product of water in pure water.

33. Since H⁺ and OH¯ are present in equal amount. [H⁺][OH¯] = 1.0× 10⁻7 M pH is employed to express these ions conc. in a convenient form. p = negative logarithm to the base 10. pH = negative logarithm to the base 10 [H⁺] pH = -Log10 [H⁺] = Log 1/[H⁺] Similarly pOH = log10[OH¯] = log 1/[OH¯] Log10 [H⁺] + log10[OH¯] = log -14 pH + pOH = 14

34. The pH value of 7 for pure water at 25°C is considered to be neutral values below 7 are acidic & above 7 basic. H⁺ pH [OH⁻]M 1 0 10¯14 0.1 1 10¯13 10¯² 2 10¯12 10¯7 7 10¯7 10¯8 8 10¯6 10¯14 14 1 10¯15 15 10

35. Clinical abnormalities of blood pH • The pH of blood may vary because of respiratory or metabolic disorders. • If pH falls below 7.4, the condition is called acidiosis. • If pH of blood rises above 7.4, the condition is called Alkalosis. • A pH imbalance caused by a change in CO₂ level is called Respiratory acidiosis or Respiratory alkalosis.

36. A pH imbalance is caused by a change in [HCO⁻₃]is called metabolic acidiosis or metabolic alkalosis. • The body compensate for a pH imbalance by adjusting the activities of the lungs or kidneys.

37. ACID BASES • Most biologic substances contain functional residue, that can accept or donate protons • So these substances are either acids or bases. • An acid is, that donates proton & base is a substance that accepts protons • Biological fluids such as blood and intracellular fluids contain a number of substances which are called as acids or bases & H+ constantly enters into the fluids. • In order to maintain a constant pH. Various biological fluids contain buffering system that serves to blunt the effect of H+

38. THE HENDERSON- HASSELBALCH EQUATION The Henderson- Hasselbalch equation is derived below. A weak acid, HA, ionizes are follows: HA H⁺ + A¯ The equilbrium constant for the dissociation is Ka= [H⁺][A¯] (H A) Cross- multiplication gives [H⁺][A¯] = Ka[HA]

39. Divide both sides by [A¯ ] : [H⁺]= Ka [HA] [A¯] Take the log of both sides log [H⁺]= log Ka[HA] [A-] log [H⁺] = log Ka + log[ HA] [A¯] Multiply through by -1: -log [H⁺]= -log Ka- log [HA] [A¯]

40. Substitute pH and pKafor –log [H⁺] and –log Ka Respectively then: pH= pKa– log [HA] [A¯] Inversion of the last term removes the minus sign and gives the Handerson- Hasselbalch equation pH= pKa + log [A¯ ] [HA]

41. The Handerson- Hasselbalch equation has great predictive value in protonicequilibria.

42. STRENTH OF ACIDS • Strong acids are completely ionized into H+ & their conjugate bases. • Weak acids dissociate to a smaller extent & the rest is remain undissociated HA H+ + A- Acid Hydrogen ion + conjugate base The dissociation constant Ka • According to low of mass action Ka = [H+] [A-] [HA] This basic expression may be written in terms of logrithms (Henderson – Hassalbalch equation) Log Ka = log [H+] + log [A-] – log [HA] By rearranging - log [H+] = - log Ka + log [A-] [HA] pH = p Ka + log [A-] [HA]

43. This equation indicates that the pH of a system will depend on its pKa, which is a constant, and the ratio of conjugate base to acid. • If the ratio of [A-] = 1 [HA] • Their pH = pKa • So pKa is that pH at which the concentration of undissociated acid (protonated) and its unprotonated forms are equal.

44. The situation where pH = pK. can never be achieved with weak acids or weak bases • In case of weak acid the conc of HA will always exceeds A- (i.e . pH <pK) and in the case of weak base pH is greater then pK. • A pH = pK can only be achieved for a weak acid by the addition of exogenous A- in the form of salt. Such as Na A. • The mixture of HA and A- is then called a buffer

45. Effect of salts upon the dissociation • when a salt of a weak acid is mixed with the weak acid in solution. • The dissociation of the acid is decreased & shows less acidity and higher pH. • Salt of weak acids are completely ionized e.g in a solution of acetic acid and sodium acetate CH3CooH CH3 Coo- +H+ CH3CooNa CH3 Coo- + Na+

46. So in this way the number of acetate ions are increased. • At the point of equilibrium, the equilibrium is shifted to the left CH3CooH CH3 Coo- + H+ • This will decrease the concentration of hydrogen ion & it also suppress the dissociation of acid & so it will resist the change of pH • Solutions containing both weak acids and their salts are referred to as buffer solutions. • They have the capacity of resisting changes of pH when weak acids or alkalis are added.

47. BUFFERS • Buffers are defined as solutions that resist changes in pH upon the addition of acids or bases. • They are the mixture of weak acids & their conjugate base, or a weak base & its conjugate acids. • A Buffer has the greatest capacity to resist such changes when pH = pK. • The Henderson – Hassebalch equation is useful in defining buffer systems in terms of pH values & concentration of buffering species.

48. Buffers of Blood Plasma • HCO3 / H2CO3 • HP-o4- / H2Po4- • Proteinate- / Protein • The bicarbonate – carbonic acid buffer system is present in greatest concentration, is important in maintaining the pH of blood. • The ratio HCO-3 / H2 Co3 in blood plasma is 20/ 1

49. So pH = pKa + log [HCO-3] [H2CO3] pH = 6.1 + log 20 – log 1 pH = 6.1 + 1.30 – 0 pH = 7.4 So the pH of the blood is 7.4

50. Buffers of red blood cells • The red cells are buffered due to the presence of protein, haemoglobin. • This is an amphoteric substance, and at the pH of red cells (7.25) it exists partly as salts. • The H Co-3 / H2 Co3 contribute to the buffer capacity of cell. • Phosphate buffer in the cell is relatively insignificant. • The H Co-3 / H2 Co3 buffer system is efficient in buffering the acids produced in the body, e.g. phosphoric, lactic, acetoacetic, ß – hydroxybutyric acids etc.