Acid-Base Balance – Anatomy, Mechanism, Functions

pH, Buffers, Acids, and Bases

Henderson–Hasselbalch equation

The Henderson–Hasselbalch equation when applied to the carbonic acid-bicarbonate buffer system in the extracellular fluids, states that:^[rx]

{\displaystyle \mathrm {pH} =\mathrm {p} K_{\mathrm {a} ~\mathrm {H} _{2}\mathrm {CO} _{3}}+\log _{10}\left({\frac {[\mathrm {HCO} _{3}^{-}]}{[\mathrm {H} _{2}\mathrm {CO} _{3}]}}\right),}

where:

• pH is the negative logarithm (or cologarithm) of molar concentration of hydrogen ions in the ECF.
• pK_{a H2CO3} is the cologarithm of the acid dissociation constant of carbonic acid. It is equal to 6.1.
• [HCO⁻
  ₃] is the molar concentration of bicarbonate in the blood plasma
• [H₂CO₃] is the molar concentration of carbonic acid in the ECF.

However, since the carbonic acid concentration is directly proportional to the partial pressure of carbon dioxide ({\displaystyle P_{{\mathrm {CO} }_{2}}}) in the extracellular fluid, the equation can be rewritten as follows:^[rx][rx]{\displaystyle \mathrm {pH} =6.1+\log _{10}\left({\frac {[\mathrm {HCO} _{3}^{-}]}{0.0307\times P_{\mathrm {CO} _{2}}}}\right),}

where:

• pH is the negative logarithm of molar concentration of hydrogen ions in the ECF.
• [HCO⁻
  ₃] is the molar concentration of bicarbonate in the plasma
• P_CO2 is the partial pressure of carbon dioxide in the blood plasma.

The pH of the extracellular fluids can thus be controlled by the regulation of PCO₂ and other metabolic acids.

Acids dissociate into H⁺ and lower pH, while bases dissociate into OH⁻ and raise pH; buffers can absorb these excess ions to maintain pH.

Self-Ionization of Water

Hydrogen ions are spontaneously generated in pure water by the dissociation (ionization) of a small percentage of water molecules into equal numbers of hydrogen (H⁺) ions and hydroxide (OH⁻) ions. The hydroxide ions remain in solution because of their hydrogen bonds with other water molecules; the hydrogen ions, consisting of naked protons, are immediately attracted to un-ionized water molecules and form hydronium ions (H₃0⁺). By convention, scientists refer to hydrogen ions and their concentration as if they were free in this state in liquid water.

$2{\text{H}}_2\text{O}\leftrightharpoons {\text{H}}_3{\text{O}}^{+}+{\text{OH}}^{-}$

The concentration of hydrogen ions dissociating from pure water is 1 × 10⁻⁷ moles H⁺ ions per liter of water. The pH is calculated as the negative of the base 10 logarithm of this concentration:

pH = −log[H⁺]

The negative log of 1 × 10⁻⁷ is equal to 7.0, which is also known as neutral pH. Human cells and blood each maintain near-neutral pH.

pH Scale

The pH of a solution indicates its acidity or basicity (alkalinity). The pH scale is an inverse logarithm that ranges from 0 to 14: anything below 7.0 (ranging from 0.0 to 6.9) is acidic, and anything above 7.0 (from 7.1 to 14.0) is basic (or alkaline ). Extremes in pH in either direction from 7.0 are usually considered inhospitable to life. The pH in cells (6.8) and the blood (7.4) are both very close to neutral, whereas the environment in the stomach is highly acidic, with a pH of 1 to 2.

Non-neutral pH readings result from dissolving acids or bases in water. Using the negative logarithm to generate positive integers, high concentrations of hydrogen ions yield a low pH, and low concentrations a high pH.

An acid is a substance that increases the concentration of hydrogen ions (H⁺) in a solution, usually by dissociating one of its hydrogen atoms. A base provides either hydroxide ions (OH⁻) or other negatively-charged ions that react with hydrogen ions in solution, thereby reducing the concentration of H⁺ and raising the pH.

Strong Acids and Strong Bases

The stronger the acid, the more readily it donates H⁺. For example, hydrochloric acid (HCl) is highly acidic and completely dissociates into hydrogen and chloride ions, whereas the acids in tomato juice or vinegar do not completely dissociate and are considered weak acids; conversely, strong bases readily donate OH⁻ and/or react with hydrogen ions. Sodium hydroxide (NaOH) and many household cleaners are highly basic and give up OH⁻ rapidly when placed in water; the OH⁻ ions react with H⁺ in solution, creating new water molecules and lowering the amount of free H⁺ in the system, thereby raising the overall pH. An example of a weak basic solution is seawater, which has a pH near 8.0, close enough to neutral that well-adapted marine organisms thrive in this alkaline environment.

Buffers

How can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers usually consist of a weak acid and its conjugate base; this enables them to readily absorb excess H⁺ or OH^–, keeping the system’s pH within a narrow range.

Maintaining a constant blood pH is critical to a person’s well-being. The buffer that maintains the pH of human blood involves carbonic acid (H₂CO₃), bicarbonate ion (HCO₃^–), and carbon dioxide (CO₂). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, hydrogen ions are removed, moderating pH changes. Similarly, excess carbonic acid can be converted into carbon dioxide gas and exhaled through the lungs; this prevents too many free hydrogen ions from building up in the blood and dangerously reducing its pH; likewise, if too much OH^– is introduced into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to jeopardize survival.

Antacids, which combat excess stomach acid, are another example of buffers. Many over-the-counter medications work similarly to blood buffers, often with at least one ion (usually carbonate) capable of absorbing hydrogen and moderating pH, bringing relief to those that suffer “heartburn” from stomach acid after eating.

Chemical Buffer Systems

Chemical buffers, such as bicarbonate and ammonia, help keep the blood’s pH in the narrow range that is compatible with life.

Acid-Base Homeostasis

Acid-base homeostasis concerns the proper balance between acids and bases; it is also called body pH. The body is very sensitive to its pH level, so strong mechanisms exist to maintain it. Outside an acceptable range of pH, proteins are denatured and digested, enzymes lose their ability to function, and death may occur.

Buffer Solution

A buffer solution is an aqueous solution of a weak acid and its conjugate base, or a weak base and its conjugate acid. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications.

Many life forms thrive only in a relatively small pH range, so they utilize a buffer solution to maintain a constant pH. One example of a buffer solution found in nature is blood. The body’s acid-base balance is normally tightly regulated, keeping the arterial blood pH between 7.38 and 7.42.

Several buffering agents that reversibly bind hydrogen ions and impede any change in pH exist. Extracellular buffers include bicarbonate and ammonia, whereas proteins and phosphates act as intracellular buffers.

The bicarbonate buffering system is especially key, as carbon dioxide (CO₂) can be shifted through carbonic acid (H₂CO₃) to hydrogen ions and bicarbonate (HCO³⁻):

${\text{H}}_2\text{O}+{\text{CO}}_2\leftrightharpoons {\text{H}}_2{\text{CO}}_3\leftrightharpoons {\text{H}}^{+}+{\text{CO}}_3^{-}$

Acid–base imbalances that overcome the buffer system can be compensated in the short term by changing the rate of ventilation. This alters the concentration of carbon dioxide in the blood and shifts the above reaction according to Le Chatelier’s principle, which in turn alters the pH.

Renal Physiology

The kidneys are slower to compensate, but renal physiology has several powerful mechanisms to control pH by the excretion of excess acid or base. In response to acidosis, the tubular cells reabsorb more bicarbonate from the tubular fluid, and the collecting duct cells secrete more hydrogen and generate more bicarbonate, and ammonia genesis leads to an increase of the NH₃ buffer.

In its responses to alkalosis, the kidneys may excrete more bicarbonate by decreasing hydrogen ion secretion from the tubular epithelial cells, and lower the rates of glutamine metabolism and ammonium excretion.

Regulation of H+ by the Lungs

Acid-base imbalances in the blood’s pH can be altered by changes in breathing to expel more CO₂ and raise pH back to normal.

Acid-base imbalance occurs when a significant insult causes the blood pH to shift out of its normal range (7.35 to 7.45). An excess of acid in the blood is called acidemia and an excess of base is called alkalemia.

The process that causes the imbalance is classified based on the etiology of the disturbance (respiratory or metabolic) and the direction of change in pH ( acidosis or alkalosis). There are four basic processes and one or a combination may occur at any given time.

• Metabolic acidosis
• Respiratory acidosis
• Metabolic alkalosis
• Respiratory alkalosis

Blood carries oxygen, carbon dioxide, and hydrogen ions (H+) between tissues and the lungs. The majority of CO₂ transported in the blood is dissolved in plasma (60% is dissolved bicarbonate).

Expiration: When blood pH drops too low, the body compensates by increasing breathing to expel more carbon dioxide.

The Role of the Kidneys in Acid-Base Balance

The kidneys help maintain the acid–base balance by excreting hydrogen ions into the urine and reabsorbing bicarbonate from the urine.

Within the human body, fluids such as blood must be maintained within the narrow range of 7.35 to 7.45, making it slightly alkaline. Outside that range, pH becomes incompatible with life; proteins are denatured and digested, enzymes lose their ability to function, and the body is unable to sustain itself.

To maintain this narrow range of pH the body has a powerful buffering system. Acid-base imbalances that overcome this system are compensated in the short term by changing the rate of ventilation.

Kidneys and Acid-Base Balance

The kidneys have two very important roles in maintaining the acid-base balance:

1. They reabsorb bicarbonate from urine.
2. They excrete hydrogen ions into urine.

The kidneys are slower to compensate than the lungs, but renal physiology has several powerful mechanisms to control pH by the excretion of excess acid or base. The major, homeostatic control point for maintaining a stable pH balance is renal excretion.

Bicarbonate (HCO₃₋) does not have a transporter, so its reabsorption involves a series of reactions in the tubule lumen and tubular epithelium. In response to acidosis, the tubular cells reabsorb more bicarbonate from the tubular fluid, and the collecting duct cells secrete more hydrogen and generate more bicarbonate, and ammonia genesis leads to an increase in the formation of the NH₃ buffer.

In response to alkalosis, the kidneys may excrete more bicarbonate by decreasing hydrogen ion secretion from the tubular epithelial cells, and lowering the rates of glutamine metabolism and ammonium excretion.

References