5. Physiology Chapter on Acid-Base Balance

The process by which the body maintains a stable pH, typically around 7.4, despite the constant production of acids.

The primary form of acid produced by the body, generated as a byproduct of energy production in cells.

The process by which the body compensates for changes in pH by altering the rate of breathing.

Acids that are not easily eliminated by the respiratory system, such as lactic acid and ketones.

The conversion of CO2 and water into carbonic acid (H2CO3).

The ability of the blood to absorb H+ ions without significant changes in pH.

The process by which the body uses its kidneys to regulate pH by eliminating excess acids or retaining bicarbonate.

The pH of arterial blood, which is slightly alkaline, typically around 7.4.

The respiratory center in the brain controls the removal of carbon dioxide (CO2) from the body, which indirectly regulates the amount of carbonic acid (H2CO3) in the extracellular fluid.

The first two lines of defense in acid-base balance are the buffer systems and the respiratory system. They work quickly to prevent major changes in pH.

The kidneys are the third line of defense in acid-base balance. They eliminate excess acid or base from the body and are the most powerful regulators in the long run.

The kidneys are slower to respond than the buffer systems and the respiratory system, taking hours to days to adjust to changes in pH.

The lungs eliminate carbon dioxide, which is a volatile acid. The kidneys remove non-volatile acids and bases.

The body produces acids from metabolism, diet, and other sources. The kidneys are responsible for excreting the net amount of acid produced.

Renal Net Acid Excretion (RNAE) is the amount of acid the kidneys excrete. It should be equal to the body's NEAP to maintain proper acid-base balance.

Acid-base balance is a complex process that involves multiple systems working together. The lungs and kidneys are key players in this delicate balance.

The phosphate buffer system plays a crucial role in regulating pH within kidney tubules, where phosphate concentration is high and pH is lower than in extracellular fluid.

Plasma proteins, particularly albumin, act as buffers in the extracellular fluid, binding both hydrogen ions (H+) and calcium ions (Ca2+).

When blood becomes acidic (acidosis), more H+ binds to plasma proteins, displacing Ca2+ and increasing free calcium levels.

In alkaline conditions (alkalemia), less H+ binds to plasma proteins, leading to more Ca2+ binding and a decrease in free calcium (hypocalcemia).

Intracellular buffers, primarily organic phosphates and proteins, maintain pH balance within cells.

Red blood cells have a rapid buffer system where hemoglobin (Hb) binds to hydrogen ions (H+) to form HHb.

While intracellular buffers are abundant, their ability to compensate for extracellular acid-base changes is usually delayed because H+ and bicarbonate (HCO3-) diffusion through cell membranes is slow.

The phosphate buffer system is less important as an extracellular buffer but plays a significant role within the kidney tubules due to high phosphate concentration and low pH.

The Henderson-Hasselbalch Equation is a mathematical formula that calculates the pH of a buffered solution. It considers the equilibrium between the acid form (HA) and the base form (A-) of the buffer and their respective concentrations.

A buffer is a solution that resists changes in pH upon addition of acid or base. It consists of a weak acid and its conjugate base working together to maintain a stable pH.

The pK value is a specific constant for each buffer pair. It reflects the strength of the weak acid in the buffer system. Stronger acids have lower pK values, indicating greater dissociation into H+ and A-. Weaker acids have higher pK values.

HA represents the acidic form of the buffer molecule, while A- represents its conjugate base form. The equilibrium constant (K) determines the ratio of these forms, influencing the pH of the solution.

The concentration of the base form of the buffer (A-) divided by the concentration of the acid form of the buffer (HA) influences the pH of the solution. A higher ratio of A- to HA will result in a higher pH, while a lower ratio will lead to a lower pH.

Distilled water lacks buffers, meaning it has no mechanism to resist changes in pH. This makes it highly susceptible to fluctuations in acidity or alkalinity.

Equilibrium constant (K) is a key factor in determining the pK value. It represents the ratio of the rates of the forward and reverse reactions in the buffer system.

The concentration of hydrogen ions (H+) defines the pH of a solution. It determines the level of acidity, with a higher concentration indicating greater acidity.

The most important buffer in the extracellular fluid, it is a combination of bicarbonate ions (HCO3-) and carbon dioxide (CO2).

The pH of the extracellular fluid is affected by the balance of bicarbonate ions (HCO3-) and carbon dioxide (CO2) in the blood.

When the body gains excess H+ ions, the buffer system shifts to the right, converting H+ and HCO3- into H2CO3, which then breaks down into CO2 and H2O. The CO2 is exhaled by the lungs.

When the body loses H+ ions, the buffer system shifts to the left, using CO2 from the lungs and water to form carbonic acid (H2CO3), which then dissociates into H+ and HCO3-.

Found in high concentrations in the cells lining the alveoli of the lungs and renal tubules. It speeds up the conversion of CO2 and water into carbonic acid and vice versa.

By altering the respiration rate, the lungs can directly affect the concentration of CO2 in the blood which in turn affects the pH.

The kidneys play a critical role in regulating the concentration of bicarbonate ions (HCO3-) in the blood, ultimately influencing the pH.

When the body's pH falls below 7.35, it is considered acidic and is called acidosis. This is often due to an increase in H+ ions in the blood.

A decrease in bicarbonate (HCO3-) concentration in the blood, leading to a lower pH.

An increase in bicarbonate (HCO3-) concentration in the blood, leading to a higher pH.

Hypoventilation (slow breathing) leads to CO2 retention, increasing PCO2 and decreasing pH.

Hyperventilation (fast breathing) leads to CO2 loss, decreasing PCO2 and increasing pH.

The primary compensatory mechanism for metabolic acid-base disturbances.