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Acid-Base Balance
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Acid-Base



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
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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]
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APPLICATIONS OF HH EQUATION

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 + HPO42-
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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-.


When urine is acidic, HCO3- combines with H+ to
form H2C03-, which is catalyzed by CA located in
the apical cell membrane of PCT.



Reabsorption is indirect.
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+
H2PO4NH3 + 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.
2.
Note whether the pH is low (acidosis) or
high (alkalosis)
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 ]) =


144 - 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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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.
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