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Acid base balance KUB by Dr. Samreena

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Acid base balance KUB by Dr. Samreena

Acid base balance KUB by Dr. Samreena

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  • 1. Acid-Base Balance
  • 2.
    • Acids are H + donors.
    • Bases are H + acceptors, or give up OH - in solution.
    • Acids and bases can be:
      • Strong – dissociate completely in solution
        • HCl, NaOH
      • Weak – dissociate only partially in solution
        • Lactic acid, carbonic acid
    Acid-Base
  • 3. pH
  • 4.  
  • 5. Buffer Systems
    • Provide or remove H + and stabilize the pH.
    • Include weak acids that can donate H + and weak bases that can absorb H + .
    • Change in pH, after addition of acid, is less than it would be in the absence of buffer.
  • 6. Chemical Buffers
    • Act within fraction of a second
    • HCO 3 - .
    • Protein.
    • Phosphate.
  • 7. HCO 3 -
    • pk= 6.1.
    • Present in large quantities.
    • Open system.
    • Respiratory and renal systems act on this buffer system.
    • Most important ECF buffer.
  • 8. Bicarbonate buffer
  • 9. Bicarbonate buffer
  • 10. Quantitative Dynamics of the Bicarbonate Buffer System
  • 11. Bicarbonate buffer
    • Sodium Bicarbonate (NaHCO 3 ) and carbonic acid (H 2 CO 3 )
    • Maintain a 20:1 ratio : HCO 3 - : H 2 CO 3
    • HCl + NaHCO 3 ↔ H 2 CO 3 + NaCl
    • NaOH + H 2 CO 3 ↔ NaHCO 3 + H 2 O
  • 12. Henderson-Hassalbalch Equation
    • pH = pK + log [base] [acid]
    • pH = pK + log [HCO - 3 ]_ Pco 2 ẋ s
    • pKa (Numerically equal to pH at which exactly on half of the protons have been removed from that group (Each component constitute (HCO - 3 & Pco 2 )50% of the total conc. of buffer system
    • If PCO 2 is expressed in:
    • kilopascals (kPa) s=0.23
  • 13. Henderson-Hassalbalch Equation
    • 0.03 millimole of H 2 C0 3 is present for each mm Hg Pco 2 measured
    • pH = 6.1 + log [HCO 3 - ] P CO2 x0.23
  • 14.  
  • 15. APPLICATIONS OF HH EQUATION
    • Physiological control of Acid-base composition of ECF
    • Use to calculate how pH of a physiologic solution responds to changes in the concentration of a week acid and/or it’s corresponding salt form.
  • 16. Proteins
    • COOH or NH 2 .
    • Largest pool of buffers in the body.
    • pk close to plasma.
    • Albumin, globulins such as Hb.
  • 17. Protein Buffers
    • Includes hemoglobin, work in blood
    • Carboxyl group gives up H +
    • Amino Group accepts H +
    • Side chains that can buffer H + are present on 27 amino acids.
  • 18. Phosphates
    • pk. = 6.8.
    • Low [ ] in ECF, better buffer in ICF, kidneys, and bone.
  • 19. Phosphate buffer
    • Major intracellular buffer
    • H + + HPO 4 2- ↔ H 2 PO4 -
    • OH - + H 2 PO 4 - ↔ H 2 O + HPO 4 2-
  • 20. Urinary Buffers
    • Nephron cannot produce a urine pH < 4.5.
    • IN order to excrete more H + , the acid must be buffered.
    • H + secreted into the urine tubule and combines with HPO 4 -2 or NH 3 .
    • HPO 4 -2 + H + H 2 PO 4 -2
    • NH 3 + H + NH 4 +
  • 21. Renal Acid-Base Regulation
    • Kidneys help regulate blood pH by excreting H + and reabsorbing HC0 3 - .
    • Most of the H + secretion occurs across the walls of the PCT in exchange for Na + .
      • Antiport mechanism.
        • Moves Na + and H + in opposite directions.
    • Normal urine normally is slightly acidic because the kidneys reabsorb almost all HC0 3 - and excrete H + .
      • Returns blood pH back to normal range.
  • 22.  
  • 23. Reabsorption of HCO 3 -
    • Apical membranes of tubule cells are impermeable to HCO 3 - .
      • Reabsorption is indirect.
    • When urine is acidic, HCO 3 - combines with H + to form H 2 C0 3 - , which is catalyzed by ca located in the apical cell membrane of PCT.
      • As [C0 2 ] increases in the filtrate, C0 2 diffuses into tubule cell and forms H 2 C0 3 .
      • H 2 C0 3 dissociates to HCO 3 - and H + .
    • HCO 3 - generated within tubule cell diffuses into peritubular capillary.
  • 24.  
  • 25. Urinary Buffers
    • Nephron cannot produce a urine pH < 4.5.
    • In order to excrete more H + , the acid must be buffered.
    • H + secreted into the urine tubule and combines with HPO 4 -2 or NH 3 .
    • HPO 4 -2 + H + H 2 PO 4 -
    • NH 3 + H + NH 4 +
  • 26.  
  • 27.  
  • 28.  
  • 29. Metabolic Acidosis
    • Gain of fixed acid or loss of HCO 3 - .
      • Plasma HCO 3 - decreases.
          • pH decreases.
  • 30. Metabolic Alkalosis
    • Loss of fixed acid or gain of HCO 3 - .
      • Plasma HCO 3 - increases.
          • pH increases.
  • 31. Metabolic Acidosis
    • Bicarbonate deficit - Blood concentrations of Bicarbonate drop below 22mEq/L
    • Causes:
      • Loss of bicarbonate through diarrhea or renal dysfunction
      • Accumulation of acids (lactic acid or ketones)
      • Failure of kidneys to excrete H+
  • 32. Compensation for Metabolic Acidosis
    • Increased ventilation
    • Renal excretion of hydrogen ions if possible
    • K + exchanges with excess H + in ECF
    • ( H + into cells, K + out of cells)
  • 33.  
  • 34. Metabolic Alkalosis
    • Bicarbonate excess - concentration in blood is greater than 26 mEq/L
    • Causes:
      • Excess vomiting = loss of stomach acid
      • Excessive use of alkaline drugs
      • Certain diuretics
      • Endocrine disorders
      • Heavy ingestion of antacids
      • Severe dehydration
  • 35. Compensation for Metabolic Alkalosis
    • Alkalosis most commonly occurs with renal dysfunction, so can’t count on kidneys
    • Respiratory compensation difficult – hypoventilation limited by hypoxia
  • 36.  
  • 37. Diagnosis of Acid-Base Imbalances
    • Note whether the pH is low (acidosis) or high (alkalosis)
    • Decide which value, pCO 2 or HCO 3 - , is outside the normal range and could be the cause of the problem. If the cause is a change in pCO 2, the problem is respiratory. If the cause is HCO 3 - the problem is metabolic.
  • 38.
    • 3. Look at the value that doesn’t correspond to the observed pH change. If it is inside the normal range, there is no compensation occurring. If it is outside the normal range, the body is partially compensating for the problem.
  • 39. Anion Gap
    • The difference between [Na + ] and the sum of [HC0 3 - ] and [Cl - ].
        • [Na + ] – ([HC0 3 - ] - [Cl - ]) =
          • 140 - 24 - 108 = 12mEq/L
            • Normal = 8-16mE/l
    • Clinicians use the anion gap to identify the cause of metabolic acidosis.
  • 40. Anion Gap
    • Law of electroneutrality:
      • Blood plasma contains an = number of + and – charges.
    • The major cation is Na + .
      • Minor cations are K + , Ca 2+ , Mg 2+ .
    • The major anions are HC0 3 - and Cl - .
        • (Routinely measured.)
      • Minor anions include albumin, phosphate, sulfate (called unmeasured anions).
      • Organic acid anions include lactate and acetoacetate,.
  • 41. Anion Gap
    • In metabolic acidosis, the strong acid releases protons that are buffered primarily by [HC0 3 ].
      • This causes plasma [HC0 3 - ] to decrease, shrinking the [HC0 3 - ] on the ionogram.
    • Anions that remain from the strong acid, are added to the plasma.
      • If lactic acid is added, the [lactate] rises.
        • Increasing the total [unmeasured anions].
      • If HCL is added, the [Cl - ] rises.
        • Decreasing the [HC0 3 - ].
  • 42. Anion Gap in Metabolic Acidosis
    • Salicylates raise the gap to 20.
    • Renal failure raises gap to 25.
    • Diabetic ketoacidosis raises the gap to 35-40.
    • Lactic acidosis raises the gap to > 35.
    • Largest gaps are caused by ketoacidosis and lactic acidosis.
  • 43.  
  • 44. THANKS

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