Ph and buffer

Bichemistry

1. P h and Buffer Acid Base Balance Prakash Pokhrel

2. PH It is the negative log of the hydrogen ion concentration. pH = -log [H + ] Acid base balance

3. pH is a unit of measure which describes the degree of acidity or alkalinity (basic) of a solution. It is measured on a scale of 0 to 14. Low pH values correspond to high concentrations of H+ and high pH values correspond to low concentrations of H+.

4. Ph value The pH value of a substance is directly related to the ratio of the hydrogen ion and hydroxyl ion concentrations. If the H+ concentration is higher than OH- the material is acidic. If the OH- concentration is higher than H+ the material is basic. 7 is neutral, < is acidic, >7 is basic

5. The ph scale The pH scale corresponds to the concentration of hydrogen ions. For example pure water H+ ion concentration is 1 x 10^-7 M, therefore the pH would then be 7 .

6. Acid Any compound which forms H⁺ ions in solution (proton donors) eg : Carbonic acid releases H⁺ ions Base Any compound which combines with H⁺ ions in solution (proton acceptors) eg:Bicarbonate (HCO3⁻) accepts H+ ions

7. Acid–Base Balance Normal pH : 7.35-7.45 Acidosis Physiological state resulting from abnormally low plasma pH Alkalosis Physiological state resulting from abnormally high plasma pH Acidemia : plasma pH < 7.35 Alkalemia : plasma pH > 7.45

8. Measurement of ph

9. Some important indicators used in a Clinical Biochemistry Laboratory are listed below: sr ,. No. INDICATOR Ph range Colour in acidic ph Colour in basic ph 1 Phenophthalein 9.3-10.5 colourless pink 2 Methyl orange 3.1-4.6 red yellow 3 Bromophenol blue 3.0-4.6 yellow blue 4 Methyl red 4.4-6.2 Red yellow 5 Phenol red 6.8 – 8.4 yellow red 6 Litmus 4.5-8.3 red Blue

10. PH meter The pH meter is a laboratory equipment which used to measure acidity or alkalinity of a solution The pH meter measures the concentration of hydrogen ions [H + ] using an ion-sensitive electrode. It is the most reliable and convenient method for measuring ph.

11. BUFFERS

12. BUFFER A buffer solution is a solution which resists changes in pH when a small amount of acid or base is added. Typically a mixture of a weak acid and a salt of its conjugate base or weak base and a salt of its conjugate acid.

13. Types of buffers Two types : ACIDIC BUFFERS – Solution of a mixture of a weak acid and a salt of this weak acid with a strong base. E.g. CH 3 COOH + CH 3 COONa ( weak acid ) ( Salt ) BASIC BUFFERS – Solution of a mixture of a weak base and a salt of this weak base with a strong acid. e.g. NH4OH + NH4Cl ( Weak base) ( Salt)

14. how buffers work Equilibrium between acid and base. Example: ACETATE BUFFER CH 3 COOH  CH 3 COO - + H + If more H + is added to this solution, it simply shifts the equilibrium to the left, absorbing H + , so the [H + ] remains unchanged. If H + is removed (e.g. by adding OH-) then the equilibrium shifts to the right, releasing H + to keep the pH constant

15. Handerson hasselbalch equation Lawrence Joseph Henderson  wrote an equation, in 1908, describing the use of carbonic acid as a buffer solution.  Karl Albert Hasselbalch  later re-expressed that formula in logarithmic terms, resulting in the Henderson– Hasselbalch equation .

16. K a = [H + ] [A - ] [HA] take the -log on both sides The Henderson- Hasselbalch Equation derivation -log K a = -log [H + ] -log [A - ] [HA] pH = p K a + log [A - ] [HA] = p K a + log [Proton acceptor] [Proton donor] HA H + + A - p K a = pH -log [A - ] [HA] apply p(x) = -log(x) and finally solve for pH…

17. - The greater the buffer capacity the less the pH changes upon addition of H + or OH - Choose a buffer whose pK a is closest to the desired pH. pH should be within pK a ± 1

18. Buffer system in body fluids

19. ACIDS VOLATILE ACIDS: Produced by oxidative metabolism of CHO,Fat,Protein Average 15000-20000 mmol of CO₂ per day Excreted through LUNGS as CO₂ gas FIXED ACIDS (1 mEq /kg/day) Acids that do not leave solution ,once produced they remain in body fluids Until eliminated by KIDNEYS Eg : Sulfuric acid ,phosphoric acid , Organic acids Are most important fixed acids in the body Are generated during catabolism of: amino acids(oxidation of sulfhydryl gps of cystine,methionine ) Phospholipids(hydrolysis) nucleic acids

20. Response to ACID BASE challenge Buffering Compensation

21. Buffers First line of defence (> 50 – 100 mEq /day) Two most common chemical buffer groups Bicarbonate Non bicarbonate ( Hb,protein,phosphate ) Blood buffer systems act instantaneously Regulate pH by binding or releasing H⁺

22. Carbonic Acid–Bicarbonate Buffer System Carbon Dioxide Most body cells constantly generate carbon dioxide Most carbon dioxide is converted to carbonic acid, which dissociates into H + and a bicarbonate ion Prevents changes in pH caused by organic acids and fixed acids in ECF Cannot protect ECF from changes in pH that result from elevated or depressed levels of CO 2 Functions only when respiratory system and respiratory control centers are working normally Ability to buffer acids is limited by availability of bicarbonate ions

23. Acid–Base Balance The Carbonic Acid–Bicarbonate Buffer System

24. The carbonic acid hydrogencarbonate buffer system The carbonic acid-hydrogen Bicarbonate ion buffer is the most important buffer system . Carbonic acid, H 2 CO 3 , acts as the weak acid Hydrogen carbonate, HCO 3 - , acts as the conjugate base Increase in H + ( aq ) ions is removed by HCO 3 - ( aq ) The equilibrium shifts to the left and most of the H + ( aq ) ions are removed

25. The small concentration of H + ( aq ) ions reacts with the OH - ( aq ) ions H 2 CO 3 dissociates, shifting the equilibrium to the right, restoring most of the H + ( aq ) ions Any increase in OH - ( aq ) ions is removed by H 2 CO 3

26. 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 Is 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

27. Phosphate Buffer System Consists of anion H 2 PO 4 - (a weak acid)(pKa-6.8) Works like the carbonic acid–bicarbonate buffer system Is important in buffering pH of ICF Limitations of Buffer Systems Provide only temporary solution to acid–base imbalance Do not eliminate H + ions Supply of buffer molecules is limited

28. Respiratory Acid-Base Control Mechanisms When chemical buffers alone cannot prevent changes in blood pH, the respiratory system is the second line of defense against changes. Eliminate or Retain CO₂ Change in pH are RAPID Occuring within minutes PCO₂ ∞ VCO₂/VA

29. 29 Phosphate buffer system The phosphate buffer system (HPO 4 2- /H 2 PO 4 - ) plays a role in plasma and erythrocytes. H 2 PO 4 - + H 2 O ↔ H 3 O + + HPO 4 2- Any acid reacts with monohydrogen phosphate to form dihydrogen phosphate dihydrogen phosphate monohydrogen phosphate H 2 PO 4 - + H 2 O ← HPO 4 2- + H 3 O + The base is neutralized by dihydrogen phosphate dihydrogen phosphate monohydrogen phosphate H 2 PO 4 - + OH - → HPO 4 2- + H 3 O +

30. Renal Acid-Base Control Mechanisms The kidneys are the third line of defence against wide changes in body fluid pH. movement of bicarbonate Retention/Excretion of acids Generating additional buffers Long term regulator of ACID – BASE balance May take hours to days for correction

31. Renal regulation of acid base balance Role of kidneys is preservation of body’s bicarbonate stores. Accomplished by: Reabsorption of 99.9% of filtered bicarbonate Regeneration of titrated bicarbonate by excretion of: Titratable acidity (mainly phosphate) Ammonium salts

32. 32 Proteins as a buffer Proteins contain – COO - groups, which, like acetate ions (CH 3 COO - ), can act as proton acceptors. Proteins also contain – NH 3 + groups, which, like ammonium ions (NH 4 + ), can donate protons. If acid comes into blood, hydronium ions can be neutralized by the – COO - groups - COO - + H 3 O + → - COOH + H 2 O If base is added, it can be neutralized by the – NH 3 + groups - NH 3 + + OH - → - NH 2 + H 2 O

33. Titratable acidity Occurs when secreted H + encounter & titrate phosphate in tubular fluid Refers to amount of strong base needed to titrate urine back to pH 7.4 40% (15-30 mEq ) of daily fixed acid load Relatively constant (not highly adaptable)

34. Renal reabsorption of bicarbonate Proximal tubule: 70-90% Loop of Henle : 10-20% Distal tubule and collecting ducts: 4-7%

35. Factors affecting renal bicarbonate reabsorption Filtered load of bicarbonate Prolonged changes in pCO2 Extracellular fluid volume Plasma chloride concentration Plasma potassium concentration Hormones (e.g., mineralocorticoids, glucocorticoids)

36. If secreted H + ions combine with filtered bicarbonate, bicarbonate is reabsorbed If secreted H + ions combine with phosphate or ammonia, net acid excretion and generation of new bicarbonate occur

37. NET ACID EXCRETION Hydrogen Ions Are secreted into tubular fluid along Proximal convoluted tubule (PCT) Distal convoluted tubule (DCT) Collecting system

38. Ammonium excretion Occurs when secreted H + combine with NH 3 and are trapped as NH 4 + salts in tubular fluid 60% (25-50 mEq ) of daily fixed acid load Very adaptable (via glutaminase induction)

39. Ammonium excretion Large amounts of H + can be excreted without extremely low urine pH because pK a of NH 3 /NH 4 + system is very high (9.2)

40. Acid – Base Balance Disturbances Interactions among the Carbonic Acid–Bicarbonate Buffer System and Compensatory Mechanisms in the Regulation of Plasma pH.

41. Four Basic Types of Imbalance Metabolic Acidosis Metabolic Alkalosis Respiratory Acidosis Respiratory Alkalosis

42. Acid Base Disorders Disorder pH [H + ] Primary disturbance Secondary response Metabolic acidosis    [HCO 3 - ]  pCO 2 Metabolic alkalosis    [HCO 3 - ]  pCO 2 Respiratory acidosis    pCO 2  [HCO 3 - ] Respiratory alkalosis    pCO 2  [HCO 3 - ]

43. Metabolic Acidosis Primary AB disorder ↓HCO₃⁻ → ↓ pH Gain of strong acid Loss of base(HCO₃⁻)

44. CAUSES OF METABOLIC ACIDOSIS LACTIC ACIDOSIS KETOACIDOSIS Diabetic Alcoholic Starvation RENAL FAILURE (acute and chronic) TOXINS Ethylene glycol Methanol Salicylates Propylene glycol

45. Acid–Base Balance Disturbances . Responses to Metabolic Acidosis

46. Metabolic acidosis Symptoms are specific and a result of the underlying pathology Respiratory effects: Hyperventilation CVS: ↓ myocardial contractility Sympathetic over activity Resistant to catecholamines CNS: Lethargy, disorientation,stupor,muscle twitching, COMA, CN palsies Others : hyperkalemia

47. Metabolic Alkalosis ↑ pH due to ↑HCO₃⁻ or ↓acid Initiation process : ↑in serum HCO₃⁻ Excessive secretion of net daily production of fixed acids Maintenance: ↓HCO₃⁻ excretion or ↑ HCO₃⁻ reclamation Chloride depletion Pottasium depletion ECF volume depletion Magnesium depletion

48. CAUSES OF METABOLIC ALKALOSIS I. Exogenous HCO3 − loads A. Acute alkali administration B. Milk-alkali syndrome II. Gastrointestinal origin 1. Vomiting 2. Gastric aspiration III. Renal origin 1. Diuretics 2. Posthypercapnic state 3. Hypercalcemia / hypoparathyroidism 4. Recovery from lactic acidosis or ketoacidosis 5. Nonreabsorbable anions including penicillin, carbenicillin 6. Mg2+ deficiency 7. K+ depletion

49. Compensation for Metabolic Alkalosis Respiratory compensation: HYPOVENTILATION ↑PCO ₂= 0.6 mm  pCO 2 per 1.0 mEq /L ↑HCO 3 - Maximal compensation: PCO ₂ 55 – 60 mmHg Hypoventilation not always found due to Hyperventilation due to pain due to pulmonary congestion due to hypoxemia(PO ₂ < 50mmHg)

50. Acid – Base Balance Disturbances . Metabolic Alkalosis

51. Metabolic Alkalosis Decreased myocardial contractility Arrythmias ↓ cerebral blood flow Confusion Mental obtundation Neuromuscular excitability

52. Respiratory Acidosis ↑ PCO ₂ → ↓pH Acute(< 24 hours) Chronic(>24 hours)

53. Compensation in Respiratory Acidosis Acute resp.acidosis : Mainly due to intracellular buffering( Hb,Pr,PO ₄) HCO₃⁻ ↑ = 1mmol for every 10 mmHg ↑ PCO ₂ Minimal increase in HCO₃⁻ pH change = 0.008 x (40 - PaCO ₂) Chronic resp.acidosis Renal compensation (acidification of urine & bicarbonate retention) comes into action HCO₃⁻ ↑= 3.5 mmol for every 10 mm Hg ↑PCO₂ pH change = 0.003 x (40 - PaCO ₂) Maximal response : 3 - 4 days

54. Acid – Base Balance Disturbances Respiratory Acid–Base Regulation.

55. Acid–Base Balance Disturbances Respiratory Acid–Base Regulation.

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