acid base balance kub by dr. samreena
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Acid base balance KUB by Dr. SamreenaTRANSCRIPT
Acid-Base Balance
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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
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pH
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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.
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Chemical Buffers
Act within fraction of a second
HCO3-.
Protein.
Phosphate.
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HCO3-
pk= 6.1.
Present in large quantities.
Open system.
Respiratory and renal systems act on this buffer
system.
Most important ECF buffer.
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Bicarbonate buffer
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Bicarbonate buffer
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Quantitative Dynamics of the Bicarbonate Buffer System
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Bicarbonate buffer Sodium Bicarbonate (NaHCO3) and carbonic
acid (H2CO3)
Maintain a 20:1 ratio : HCO3- : H2CO3
HCl + NaHCO3 ↔ H2CO3 + NaCl
NaOH + H2CO3 ↔ NaHCO3 + H2O
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Henderson-Hassalbalch Equation
pH = pK + log [base] [acid]
pH = pK + log [HCO-3]_
Pco2 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 & Pco2 )50% of the total conc. of buffer system
If PCO 2 is expressed in:
kilopascals (kPa) s=0.23
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Henderson-Hassalbalch Equation
0.03 millimole of H2C03 is present for
each mm Hg Pco2 measured
pH = 6.1 + log [HCO3-]
PCO2 x0.23
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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.
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Proteins
COOH or NH2.
Largest pool of buffers in the body.
pk close to plasma.
Albumin, globulins such as Hb.
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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.
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Phosphates
pk. = 6.8.
Low [ ] in ECF, better buffer in ICF,
kidneys, and bone.
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Phosphate buffer
Major intracellular buffer
H+ + HPO42- ↔ H2PO4-
OH- + H2PO4- ↔ H2O + HPO4
2-
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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 HPO4-2 or NH3.
HPO4-2 + H+ H2PO4
-2 NH3 + H+ NH4
+
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Renal Acid-Base Regulation
Kidneys help regulate blood pH by excreting H+ and reabsorbing HC03
-. 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 HC03
- and excrete H+. Returns blood pH back to normal range.
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Reabsorption of HCO3-
Apical membranes of tubule cells are impermeable to HCO3
-. Reabsorption is indirect.
When urine is acidic, HCO3- combines with H+ to
form H2C03-, which is catalyzed by ca located in the
apical cell membrane of PCT. As [C02] increases in the filtrate, C02 diffuses into tubule
cell and forms H2C03. H2C03 dissociates to HCO3
- and H+. HCO3
- generated within tubule cell diffuses into peritubular capillary.
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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 HPO4-2 or NH3.
HPO4-2 + H+ H2PO4
-
NH3 + H+ NH4+
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Metabolic Acidosis
Gain of fixed acid or loss of HCO3-.
Plasma HCO3- decreases.
pH decreases.
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Metabolic Alkalosis
Loss of fixed acid or gain of HCO3-.
Plasma HCO3- increases.
pH increases.
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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+
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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)
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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
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Compensation for Metabolic Alkalosis
Alkalosis most commonly occurs with renal
dysfunction, so can’t count on kidneys
Respiratory compensation difficult –
hypoventilation limited by hypoxia
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Diagnosis of Acid-Base Imbalances
1. Note whether the pH is low (acidosis) or high (alkalosis)
2. Decide which value, pCO2 or HCO3- , is
outside the normal range and could be the cause of the problem. If the cause is a change in pCO2, the problem is respiratory. If the cause is HCO3
- the problem is metabolic.
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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.
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Anion Gap
The difference between [Na+] and the sum of [HC03
-] and [Cl-].
[Na+] – ([HC03-] - [Cl-]) =
140 - 24 - 108 = 12mEq/L Normal = 8-16mE/l
Clinicians use the anion gap to identify the cause of metabolic acidosis.
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Anion Gap
Law of electroneutrality: Blood plasma contains an =
number of + and – charges. The major cation is Na+.
Minor cations are K+, Ca2+ , Mg2+.
The major anions are HC03- and
Cl-. (Routinely measured.)
Minor anions include albumin, phosphate, sulfate (called unmeasured anions).
Organic acid anions include lactate and acetoacetate,.
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Anion Gap
In metabolic acidosis, the strong acid releases protons that are buffered primarily by [HC03].
This causes plasma [HC03-] to decrease,
shrinking the [HC03-] 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 [HC03-].
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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.
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