Acid-Base Balance

Acid-Base Balance

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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


pH



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.


Chemical Buffers

Act within fraction of a second

HCO3-.

Protein.

Phosphate.


HCO3-

pk= 6.1.

Present in large quantities.

Open system.

Respiratory and renal systems act on this buffer

system.

Most important ECF buffer.


Bicarbonate buffer


Bicarbonate buffer


Quantitative Dynamics of the Bicarbonate Buffer System


Bicarbonate buffer Sodium Bicarbonate (NaHCO3) and carbonic

acid (H2CO3)

Maintain a 20:1 ratio : HCO3- : H2CO3

HCl + NaHCO3 ↔ H2CO3 + NaCl

NaOH + H2CO3 ↔ NaHCO3 + H2O


Henderson-Hassalbalch Equation

pH = pK + log [base] [acid]



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.


Proteins

COOH or NH2.

Largest pool of buffers in the body.

pk close to plasma.

Albumin, globulins such as Hb.


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.


Phosphates

pk. = 6.8.

Low [ ] in ECF, better buffer in ICF,

kidneys, and bone.


Phosphate buffer

Major intracellular buffer

H+ + HPO42- ↔ H2PO4-

OH- + H2PO4- ↔ H2O + HPO4

2-


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

+


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.



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.



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+





Metabolic Acidosis

Gain of fixed acid or loss of HCO3-.

Plasma HCO3- decreases.

pH decreases.


Metabolic Alkalosis

Loss of fixed acid or gain of HCO3-.

Plasma HCO3- increases.

pH increases.


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+


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)



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


Compensation for Metabolic Alkalosis

Alkalosis most commonly occurs with renal

dysfunction, so can’t count on kidneys

Respiratory compensation difficult –

hypoventilation limited by hypoxia



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.


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.


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.


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,.


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-].

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.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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Acid-Base Balance

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