While long term disturbances in the acid–base balance can disturb the normal growth and development, acute short term, but severe changes in pH, can also be drastic. The body's acid base balance is maintained by different organ systems, including the lungs and the kidneys, along with the intracellular and extracellular buffer systems. A human body’s normal pH ranges from 7.35-7.45. pH and hydrogen ions concentration have an inverse relationship with each other. Changes in pH can force the hydrogen ions to combine with proteins, disturbing their function.
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Acid-base reactions

Image: “Acid-Base Reactions” by Phil Schatz. License: CC BY 4.0


Review of [H+] and pH

pH = -log10 [H+]

Average arterial blood pH

  • pH = 7.4
  • [H+] = 0.0000004

Thus, large changes in pH, such as in the GI system, require 10 fold changes in hydrogen concentration.

  • pH = 1; [H+] = 0.1
  • pH = 8; [H+] = 0.00000001
ph-indicator-strip

pH Indicator Strip

Example pH values of various fluids

Compartment pH
Gastric secretions (under conditions of maximal acidity) 0.7 → H+-K+-ATPase (proton pump) in parietal cells
Lysosome 5.5
Chromaffin granule 5.5
Neutral H2O at 37°C 6.81
Cytosol of a typical cell 7.2 Range: 6.0—7.4 based in part on metabolism
Cerebrospinal fluid 7.3
Arterial blood plasma 7.4 Range: 7.35—7.45 highly regulated
Mitochondrial inner matrix 7.5
Secreted pancreatic fluid 8.1 → HCO3 secretion ductal cells

Blood pH: Acidosis and Alkalosis

normal-blood-ph

Arterial blood pH

acidemia-alkalemia

Acids and Bases

A substance that releases hydrogen ion (H+) or dissociates into a hydrogen ion and a conjugate base (A−) is termed as acid. Base is a substance that accepts a hydrogen ion. If HA is added, some concentration of HA dissociates and H+ and A− concentrations increase so that the equilibrium is maintained. Similarly, adding H+ ions joins with A- leading to a decrease in A− concentration, resulting in an increase in HA.

Buffers

Substances that try to maintain a normal range of pH when an acid or a base is added in various reactions are called buffers. In the absence of buffers, a minor change in H+ ions concentration can lead to a dramatic change in the pH. Buffers actually avoid large changes in pH by reacting with the new generated H+ ions, thus neutralizing them:

A +  H+ → HA

Weak acids and bases are known to be the best buffers because a buffer which is 50% dissociated i.e. equally into HA and A−; pH, at which this occurs, is called pK ionization constant. The best buffers for the human body are those which have a pK close to 7.40; these include the physiological buffers of the human body. If ph < pk of a buffer, HA > A− and when ph > pk, A− is more than HA.

Physiologic Buffers

These include bicarbonate and non-bicarbonate buffers which are actually protective for the body against abnormal changes in the pH. While managing patients, doctors routinely monitor the bicarbonate buffer system. In this system, bicarbonate picks up surplus H+ ions from the system, producing carbon dioxide (CO2) and water. Similarly, when H+ ion concentration decreases in the system, carbon dioxide combines with water to release hydrogen and bicarbonate ions. The pK of this reaction is 6.1. The Henderson-Hasselbalch equation expresses the relationship among pH, pk, and the concentrations of an acid and its conjugate base which is valid for any buffer. The Henderson-Hasselbalch equation for bicarbonate and carbon dioxide is as follows:

pH = 6.1 + log [HCO3]/[CO2]

The Henderson-Hasselbalch equation contains 3 variables; if any two of these are known, the third can be calculated. CO2 is taken clinically as mm Hg and multiplied by its solubility constant, 0.03 mmol/L/mm Hg.

  1. The non-bicarbonate buffers

The non-bicarbonate buffers include proteins and phosphates. Protein buffers include the extracellular proteins (albumin) and intracellular proteins (hemoglobin). Because of the amino acid histidine in the proteins, this can bind or release hydrogen ions. Proteins are very effective buffers. On the basis of its position in the protein molecule, its pK can vary; the average pK is about 6.5.

Hemoglobin and albumin both have histidine 34 and 16, respectively.

  1. Phosphate Buffers:

Phosphates can bind 3 H+. They can exist as PO43−, HPO4 2−, H2PO4 1−, or H3PO4. Mostly they exist as HPO4 2− or H2PO4 1−.

H2PO4 1-  →   H+ + HPO42-

As the pK of this reaction is 6.8, it is also a very good buffer. It is of lesser significance than albumin because concentration of PO4 in the ECF is relatively low, but it is found in a higher concentration in the urine, where it is an important buffer.

Extracellular pH and Intracellular pH

It is actually the intracellular pH that affects the cell function. Any change in the intracellular pH also changes the extracellular pH in one way or another. Although intracellular pH is more important, we only measure the extracellular pH clinically while managing patients. Still, the intracellular buffer system is more active and efficient in managing major pH changes.

Normal Acid – Base Balance

Two major organs are involved in the normal acid and base balance. Major weak acid generated in the body as a result of metabolism is CO2; this depends on the physical activity.

Lungs excrete CO2 from the system out into the environment and maintain a Pco2 in the blood. Central ventilation control in the CNS controls the rate of breathing in order to maintain Pco2 between 35-45 mmHg. When the ventilation rate increases, Pco2 decreases and when the rate of ventilation decreases, Pco2 increases.

Kidneys remove endogenous acids from the body through the urine. On average 1-2 mEq/kg/24 hr of H+ ions are produced in the body in adults, whereas, in children, 2-3 mEq/kg/24 hr of H+ ions are produced. Main sources include dietary protein metabolism, incomplete metabolism of carbohydrates and fat, and stool losses of bicarbonate. This shows that acid production depends on the amount of protein intake in the diet and the rate of catabolic activity. No H+ ions are produced after complete carbohydrate and fat metabolism; instead, water and CO2 which is released through the lungs outside. Incomplete catabolism of carbohydrates and fats produce lactic acid and keto acids, like β-hydroxybutyric acid and acetoacetic acid. Normally, this occurs in small amounts but, in various pathologic conditions, like lactic acidosis and diabetic ketoacidosis, endogenous acid production is abnormally increased. Stool loss of bicarbonate also contributes to the overall acid production in the body. The stomach mainly produces H+ ions, the rest of the GI tract produces bicarbonate with an overall loss of bicarbonate from the body into the GI tract and out through the stools. To produce bicarbonate into the GI tract, H+ ions are released into the blood stream, i.e. one H+ ion for each bicarbonate. This type of acid production is also not much in the body, but it is dramatically increased in a patient with watery stools.

Only the lungs can regulate the CO2 concentration, and only the kidneys can regulate the bicarbonate concentration.

Excess H+ ions with endogenous acid producing mechanisms in the body are neutralized by the bicarbonate, leading to a decrease in its concentration. This bicarbonate deficiency is restored by the kidneys secreting hydrogen ions. Lungs cannot regenerate bicarbonate, although the loss of carbon dioxide also lowers the hydrogen ion concentration:

H+ +  HCO3  → CO2 +H2O

During the metabolic acidosis, an increased ventilation rate lowers the CO2 concentration in the blood, moving this reaction towards the right and decreasing H+ ion concentration, increasing the pH. Metabolic acidosis still persists in the body.

Similarly, the renal mechanism cannot actually tolerate and correct an unusually raised CO2 concentration. The following reaction is occurring in the body:

H+  +   HCO3 →  CO2 +  H2O

Due to metabolic activity, raised bicarbonate concentration or more H+ ion concentration promotes the forward reaction increasing the CO2 concentration and decreasing the hydrogen ion concentration. In respiratory acidosis, the bicarbonate concentration is increased in the kidneys which buffers and decrease H+ ion concentration and maintain the pH, but the kidneys cannot remove the raised CO2 which is removed by the lungs. Both the lungs and the kidneys can affect the hydrogen ion concentration, hence the pH. Only the lungs can regulate the CO2 concentration, and only the kidneys can regulate the bicarbonate concentration.

Henderson-Hasselbalch-Examples

pH = 6.1 + log (HCO3 / (0.03 x PCO2))

Normal arterial blood gas Acidic arterial blood gas Alkalotic arterial blood gas
HCO3 = 24 mM HCO3 = 26 mM HCO3 = 22 mM
PaCO2 = 40 mmHg PaCO2 = 60 mmHg PaCO2 = 20 mmHg
pH = 7.40 pH = 7.26 pH = 7.66

Renal Handling of Acids and Bases

After reclaiming filtered bicarbonate, the renal acid–base handling mechanism allows the excretion of acid formed by endogenous acid production (explained earlier). It occurs in the collecting ducts and distal tubule. As the hydrogen pumps in collecting ducts cannot lower the urine pH below 4.5, excretion of endogenous acids need the presence of urinary buffers.

The two main urinary buffers are

  1. Phosphate buffer
  2. Ammonia buffer

The concentration of phosphates in the urine depends on the amount of dietary intake and the amount filtered and later reabsorbed in the proximal tubules. The serum level of phosphates is much lower than the concentration of urinary phosphates. These phosphates serve as an effective buffer through:

H+ +  HPO4 2−  → H2PO4 −                                                  pK = 6.8 

As the urine pH decreases from 7.0 to 5.0 within the collecting duct, phosphate buffers are very effective, but the buffering capacity of this reaction is limited by its concentration. Urinary phosphate concentration cannot be modified by any mechanism.

Renal Mechanisms in Acid-Base Balance

The major mechanism, by which kidneys maintain pH, is by regulating serum bicarbonate concentration promoting H+ excretion in the urine. This is a two-step process.

  1. Initially, tubules reabsorb bicarbonate filtered from the blood; then
  2. Tubular secretion of hydrogen ions occurs

Excretion of H+ into the urine leaves behind the HCO3 ions which neutralize the endogenous acid production. The renal acid excretion occurs throughout the nephron directly or indirectly. Urinary buffers maintain the pH and allow the excretion, or neutralization, of endogenous acids.

Bicarbonate re-absorption occurs in the proximal tubules.

Ammonia: NH3 & NH4

Ammonia production can be modified, allowing for the regulation of acid excretion. The buffering capacity of ammonia is based on the reaction:

NH3 +  H+  →     NH4+

Proximal tubular cells excrete ammonia, after the metabolism of glutamine amino acid:

Glutamine → ­NH4+ +  glutamate

Glutamine → NH4+ + ketoglutarate

One glutamine produces two NH4+ ions. Later, the metabolism of α-ketoglutarate also produces two bicarbonate ions. Ammonium ions go into the lumen of tubule, and bicarbonate ions exit the proximal tubule cells via Na+, 3 HCO3 co-transporter. Cells of the thick ascending limb of the loop of Henle re-absorb the ammonium ions which later moves into the blood neutralizing the bicarbonate ion produced earlier in the proximal tubules. These ammonium ions are actually the sources of H+ ions, making the cells of collecting ducts at a crucial place in excretion of H+ in the form of NH4+ ions.

Cells of CD secrete H+ ions and regenerate bicarbonate, which moves into the bloodstream and neutralizes the endogenous acid production. Renal buffers, phosphate and ammonia buffer the H+ secreted. Ammonia is present in a higher concentration in the renal interstitium and, due to free permeability in collecting ducts, it diffuses into the lumen and neutralizes the H+ ions responsible for the low urine pH; effectively buffering the pH. Renal H+ ion excretion is regulated according to the body’s physiologic requirements. When more acid production occurs, ammonia production is also up regulated in the proximal tubules to serve as a buffer in the collecting ducts i.e. NH3 excretion may rise up to 10-folds over the baseline values.

Extracellular pH: Regulator of Renal Acid Excretion

Extracellular pH is the most important regulator of renal acid excretion. A fall in extracellular pH with respiratory or metabolic cause promotes renal acid excretion. Through the renin-angiotensin system, aldosterone also stimulates H+ ion excretion in the CD resultant increase in serum HCO3 concentration. Hypokalemia also increases acid secretion, by promoting ammonia production in the PCT and increasing H+ ion secretion in the CD; leading to metabolic alkalosis. Opposite effects occur in hyperkalemia leading to metabolic acidosis.

In case of alkalosis, bicarbonate resorption in the PCT is decreased; losing it into the urine and secretion of H+ ions by the cells of CD into the lumen is reduced. Bicarbonate is lost in the urine and H+ ions move into the blood, compensating the body’s alkalosis.

 

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