Definition
Acid-base balance regulates the body’s hydrogen ion (H⁺) concentration to maintain the blood’s pH within the normal range of 7.35–7.45. Maintaining this balance is crucial for cellular function and overall homeostasis.
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- The normal pH of arterial blood is 7.4
- pH of venous blood and interstitial fluids is about 7.35
- The blood pH is maintained within a remarkable constant level of 7.35 to 7.45.
Important of pH
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- Changes in pH affect the ionization of protein molecules and, consequently activity of many enzymes.
- Changes in pH, together with the partial pressure of carbon dioxide (pCO2), change the shape of the hemoglobin.
- A decrease in pH increases sympathetic tone and may lead to cardiac dysrhythmias.
Metabolic sources of acids
During metabolic processes, two types of acids are produced
Fixed acids or non-volatile acids:
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- Phosphoric
- Sulphuric acids
- Pyruvic acid,
- Lactic acid
- Keto acids
Volatile acids:
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- Carbonic acid (H2CO3).
Metabolic sources of bases
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- Citrate salts of fruit juices may produce bicarbonate salt.
- Deamination of amino acids produces ammonia
- The formation of bis-phosphate also contributes to the alkalinizing effect.
Maintenance of normal blood ph
To maintain the blood pH at 7.35 –7.45, three primary systems regulate the hydrogen ion concentration in the body fluids. These are:
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- Buffer mechanism: First line of defense
- The respiratory mechanism: Second line of defense
- Renal mechanism: Third line of defense.
Buffer systems and their role in acid-base balance
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- A buffer is a mixture of a weak acid and a salt of its conjugate base.
- A buffer can reversibly bind hydrogen ions. Free H+ combines with the buffer to form a weak acid (H buffer)
Blood Buffers
Various buffer systems present in the human body are
Buffers of extracellular fluid are present in plasma.
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- Bicarbonate buffer
- Phosphate buffer
- Protein buffer
Buffers of intracellular fluid present in RBCs
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- Bicarbonate buffer
- Phosphate buffer
- Hemoglobin buffer
Bicarbonate Buffer System (HCO3– / H2CO3)
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- The bicarbonate buffer system is the most important extracellular buffer.
- It plays an important role in maintaining blood pH because of its high concentration.
- Two elements of the buffer system, bicarbonate (HCO3–) and carbonic acid (H2CO3), are regulated by the kidneys and by the lungs, respectively.
- Under physiological conditions, with a plasma pH 7.4, the bicarbonate to carbonic acid (HCO3– /H2CO3) ratio is 20:1.
Mechanism of action of bicarbonate buffer
When a strong acid, such as HCI, is added to the bicarbonate buffer solution, the increased hydrogen ions are buffered by HCO3– to form the very weak acid H2CO3, which, in turn, forms CO2 and H2O.
When sodium hydroxide (NaOH) is added to bicarbonate buffer, hydroxyl ion (OH–) from NaOH combines with H2CO3 to form weak base HCO3– and H2O
Phosphate Buffer System (HPO4– –/H2PO4–)
The phosphate buffer system is less important than a blood buffer; it plays a major role in buffering renal tubular fluid and intracellular fluids.
Mechanism of Action of Phosphate Buffer
When a strong acid such as HCI is added to phosphate buffer, the H+ is accepted by the base (hydrogen phosphate) HPO4– – and converted to (dihydrogen phosphate) H2PO4– and the strong acid HCI is replaced by a weak acid (sodium dihydrogen phosphate) NaH2PO4
When a strong base, such as NaOH, is added to the phosphate buffer, the OH– is buffered by the H2PO4– to form HPO4– – and water. Thus, strong base NaOH is replaced by weak base HPO4– –
At a plasma pH 7.4, the HPO4– – : H2PO4– is 4:1.
Protein Buffer (Na Protein/H Protein)
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- In the blood, plasma proteins, especially albumin, act as a buffer.
- In acid solution, the basic amino group (NH2) takes up excess H+ ions forming (NH3+).
- In basic solutions, the acidic (carboxylic acids) COOH groups give up hydrogen ions forming (hydroxide) OH– of alkali to water.
Hemoglobin Buffer
Hemoglobin is the major intracellular buffer of blood which is present in erythrocytes. It buffers carbonic acid (H2CO3) and anhydride CO2 from the tissues.
Action of hemoglobin buffer
Respiratory Mechanism
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- Second line of defense against acid-base disturbances
- It regulates the concentration of carbonic acid (H2CO3) in blood and other body fluids by the lungs.
- The respiratory center regulates the lungs’ removal or retention of CO2 and H2CO3 from the extracellular fluid.
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- An increase in (H+) or (H2CO3) stimulates the respiratory center to increase the rate of respiratory ventilation, and excess acid (H2CO3) in the form of CO2 is quickly removed
- An increase in (OH–) or (HCO3–) depresses respiratory ventilation and release of CO2 from the blood
- The increased blood CO2 will result in the formation of more H2CO3 acid. To neutralize excess alkali (HCO3–)
Renal mechanism in acid-base balance
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- The renal mechanism is the third line of defense in acid-base balance. Renal mechanisms exert long-term acid-base control.
- The kidney regulates acid-base balance by conserving HCO3– (alkali reserve) and acid excretion.
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- The pH of the initial glomerular filtrate is approximately 7.4, whereas the average urinary pH is approximately 6.0 due to the excretion of non-volatile acids produced by metabolic processes.
- The pH of the urine may vary from 4.5 to 8.0 corresponding to the case of acidosis or alkalosis.
- This ability to excrete variable amounts of acid or base makes the kidney the final defense mechanism against changes in body pH
Renal four key mechanisms
1.Exchange of H+ for Na+ of tubular fluid.
2. Reabsorption of bicarbonate from tubular fluid.
3. Formation of ammonia and excretion of ammonium ion (NH4+) in the urine.
4.Excretion of H+ as (dihydrogen phosphate) H2PO4– in urine
Exchange of H+ for Na+ of tubular fluid and reabsorption of bicarbonate from tubular fluid.
Formation of ammonia and excretion of ammonium Ions in the urine.
Excretion of H+ as H2PO4 – in urine.
Disorders of Acid-Base Balance
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- Acidemia is defined as an arterial blood pH of less than 7.35.
- Alkalemia is defined as an arterial blood pH of greater than 7.45.
- Acidosis and alkalosis refer to pathological states that can lead to acidemia or alkalemia.
Acidosis and alkalosis are classified in terms of their cause :
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- Metabolic acidosis: decrease in bicarbonate (HCO3– )
- Metabolic alkalosis: increase in bicarbonate (HCO3– )
- Respiratory acidosis: increase in pCO2 or H2CO3
- Respiratory alkalosis: decrease in pCO2 or H2CO3