1. Dr. Saidunnisa Professor of Biochemistry Acids, bases, conjugate acid base pairs, body buffers.

2. Learning objectives At the end of the session student shall be able to : List the major sources of acids and bases in the body. Define an acid ; base ; buffer; conjugate acid base pairs with suitable examples. List the various buffer systems. Revise your concept of pH ; pKa and Henderson – hasselbalch equation Study how the bicarbonate phosphate protein and Hemoglobin buffer system works Explain how the above are linked with respiration and kidneys. What is meant by the term Alkali reserve.

3. Introduction Under normal conditions the pH of ECF usually does not vary beyond the range 7.35- 7.45 and is maintained approximately at 7.4. Maintenance of this pH is one of the prime requisites of life and any variation on either side seriously disturbs the vital processes and may lead to death.

4. pH less than 7.3 leads to acidosis causing CNS depression, coma and death. pH more than 7.5 leads to alkalosis which induces neuromuscular hyper excitability and tetany and death. Large amounts of H + are continuously contributed to ECF from intracellular metabolic reactions. Hence to maintain a constant pH it is necessary that they are removed from the fluids promptly and effectively.

5. Productions of acids by the Body Carbonic acid (H 2 CO 3 ):- Chief acid produced in the body during oxidation of carbon compounds Sulphuric acid (H 2 So 4 ) :-Produced during oxidation of ‘S’ –containing aminoacids Phosphoric acid(H 2 Po 4 ) :- Produced during metabolism of phospho proteins and nucleoproteins Organic acids:- Pyruvic acids, Lactic acids, and ketone bodies are the intermediates in the metabolism Iatrogenic :- Antacids Diet rich in animal protein

6. Production of bases by the body Formation of basic compounds in the body in normal circumstances is negligible. Bicarbonate produced from organic acids. Diet rich in vegetables.

7. Acid/Base definitions Definition #1: Arrhenius (traditional) Acids – produce H + ions (or hydronium ions H 3 O + ) Bases – produce OH - ions (problem: some bases don’t have hydroxide ions!)

8. Acid/Base Definitions Definition #2: Brønsted – Lowry Acids – proton donor Bases – proton acceptor A “proton” is really just a hydrogen atom that has lost it’s electron!

9. ACID-BASE THEORIES The Brønsted definition means NH 3 is a BASE in water — and water is itself an ACID

10. Conjugate Pairs

11. Learning Check! Label the acid, base, conjugate acid, and conjugate base in each reaction: HCl + OH -  Cl - + H 2 O H 2 O + H 2 SO 4  HSO 4 - + H 3 O +

12. Conjugate acid-base A strong acid and a weak base and vice-versa are called conjugate acid-base pairs. Acid and Bases can be : Strong: dissociate completely HCL H + + Cl - (complete) Strong weak base Acid The concentration of H+ is very high in strong acid.

13. – Weak – dissociate only partially in solution carbonic acid, Lactic acid, H 2 CO 3 H + + HCO 3 - (partially) Weak acid The number of acid molecules existing may be only 50%.

14. The weak acids exist in two forms: The dissociated form will release H + ions and the conjugate base For e.g acetic acid, the dissociation will give rise to free H + ions and acetate ions The undissociated form existing as acetic acid. CH3COOH CH3COO- + H+

15. pKa- Definition The pH at which the rate of forward reaction (dissociation) is in equilibrium with the rate of backward reaction. It is designated as pKa or dissociation constant. CH3COOH CH3COO- + H+ Ka = CH3COO- + H+ CH3COOH (undissociated molecules) Strong acids have low pKa and weak acids have higher pKa.

16. pKa-Importance The ratio between the dissociated and un dissociated form is equal to 1 at a pH equal to its pKa. Example in the case of acetic acid, at a pH of 4.76 (pKa), the amount of dissociated form is equal to the un dissociated form. The combination of these two forms is buffer.

17. Acidity of the solution Is equal to the ratio of the activities of the acid to the base multiplied by the dissociation constant. [H+] = Ka [Acid] or [HA] [Base] [A-] Ka = Dissociation constant.

18. pH- Definition Sorensen expressed the H+ concentration as the negative logarithm of hydrogen ion activity and is designated as the pH. pH = -log [H+] pH value is inversely proportional to the acidity.

19. Henderson-Hasselbalch equation Effects of salt upon the dissociation of acids. Relationship between pH, pKa, concentration of acid and base expressed by Henderson- Hasselbalch equation. pH = pKa + log [Base] or pH = pKa + log [Salt] [Acid] [Acid] when the concentration of the base and the acid are the same then pH= pKa.

20. Indicators Weak organic acids or bases They change in colour during pH change. The change is over a specific range only. Each indicator exists as conjugate pair Each member differs sharply in colour. It follows the law of Handerson –Hasselbalch equation.

21. An average rate of metabolic activity produces roughly 22,000mEq acid per day. If all the acids were dissolved at one time in unbuffered body fluids their pH would be less than 1. However the pH of blood is maintained between 7.36-7.44. How is it possible?

22. Buffer A buffer is defined as a solution which resist the change in pH when an acid or alkali is added. HCL + NaHCO 3 H 2 CO 3 + NaCl Strong Buffer weak Neutral Acid Acid salt

23. Buffers can resist the change in pH when an acid or alkali is added upto approximately ±pKa 1.

24. Buffers can act quickly but not permanently. It cannot remove the H + ions from the body. It acts as a temporary measure to decrease the free H + ions. Final elimination is through lungs or kidneys.

25. The body has developed three mechanisms to regulate the body’s acid –base balance and maintain the ECF at pH 7.4 BUFFERS RESPIRATORY MECHANISMS RENAL MECHANISMS

26. Body buffers-Distribution 1. Bicarbonate-Carbonic acid buffer 2. Hemoglobin buffer 3. Phosphate buffer 4. Protein buffer 5. Renal cells-Ammonia buffer Relative distribution of various buffer systems in the body :- 42% are present in ECF and 58% in ICF

27. Buffer systems of the body 9.25 Renal cells-Ammonia buffer pKa 9.25

28. Bicarbonate buffer system It is most predominant buffer system of the ECF particularly the plasma H 2 CO 3 H+ HCO 3 - By the law of mass action pKa= [H+ ] [HCO 3 -] [ H 2 CO 3 ] pKa = dissociation constant of carbonic acid = 6.1

29. The equation may be rewritten as [H+ ] = pKa [H 2 CO 3 ] [HCO 3 - ] pH = log 1 [H + ] By taking the reciprocals and logarithms we get pH =pKa + log [ HCO 3 - ] = [Base] [H 2 CO 3 ] [Acid] This reaction is called as Henderson – Hassel Balch equation which is valid for any buffer pair.

30. NaHCO 3 = 20 H 2 CO 3 1 At blood pH 7.4 the ratio of bicarbonate to carbonic acid is 20:1. Bicarbonate is much higher i.e 20 times higher than carbonic acid in blood. This is referred to as alkali reserve and is responsible for effective buffering of H+ ions generated in the body.

31. Neutralization Eg:- Neutralization of strong and non-volatile acids such as HCL,H 2 SO 4 and lactic acid entering the ECF is achieved by bicarbonate buffer. Thus lactic acid is buffered as follows:

32. Carbonic Acid thus formed is volatile and is eliminated by diffusion of CO 2 through alveoli of lungs Note: Proper lung functioning is important hence bicarbonate buffer is linked up with respiration.

33. Similarly when alkaline substances e.g NaOH enters the ECF it reacts with the acid component that is H 2 CO 3 of the buffer system. Alkali reserve is represented by the sodium bicarbonate concentration in the blood that has not combined with strong and non-volatile acids.

34. Advantages It is a very good physiological buffer and acts as a first line of defense because It is present in very high concentration than other buffer systems Produces carbonic acids which is weak and volatile acid CO 2 is exhaled out. Disadvantages: As a chemical buffer it is rather weak. pKa is further away from the physiological pH.

35. Phosphate Buffer System Salt / Acid = ( Na 2 HPO 4 / NaH 2 PO 4 ) It is mostly an intracellular buffer and less important in plasma due to its low concentration with pKa of 6.8 close to blood pH 7.4. It would have been more affective had it been present in high concentration. Normal ratio in plasma is 4:1 The phosphate buffer system is directly linked up with the kidneys.

36. PH = pKa + log [ salt ] [ acid ] 7.4 = 6.8 + log [ salt ] [ Acid ] 0.6 anti log is 4

37. Phosphate buffer E.g : When strong acid enters the blood it is buffered by base of the phosphate as follows When the alkali enters the blood it is buffered by the acid of phosphate as follows

38. When the alkali enters the blood it is buffered by the acid of phosphate as follows Thus phosphate buffer system works in conjunction with the kidneys. A normal healthy kidney is necessary for proper functioning Advantages: As a chemical buffer it is very effective as pKa approaches the physiological pH. Disadvantage: As physiological buffer it is less efficient because it is present in low concentration in the Blood.

39. Protein Buffer System Salt / Acid = Na + Pr - / H + Pr -. Buffering capacity of plasma proteins much less than Hb. Eg: Hb present in one litre if blood can buffer 27.5mEq of H+ ions Where as plasma proteins can buffer 4.24mEq of H+ ions at pH 7.5 In acidic medium protein acts as a base amino group takes up H + ions from the medium forming NH 3 + and proteins become positively charged. In alkaline medium protein acts as an acid COOH group dissociates and gives H+ forming COO- so proteins become negatively charged.

40. Hb as a buffering System Buffering capacity of Hb is due to the presence of imidazole nitrogen group of histidine amionoacid. 38 histidine residues are present in 1 molecule of hb and has a pka 6.1. Oxygenated Hb is Stronger acid than deoxygenated Hb. On oxygenation; the imidazole N 2 group acts as acid and donates protons in the medium. Deoxygenated Hb ; the imidazole N 2 group acts as a base and takes up protons from the medium.

41. Acidity of the medium favours delivery of oxygen ; alkalinity favours oxygenation of Hb. The role of the Hb buffer is considered along with the respiratory regulation of pH.

42. Ammonia buffer Ammonia is an urinary buffer. Ammonia is toxic and hence not present in the blood, but generated in the renal tubules to combine with H + ions to be excreted as ammonium salts.