Acid-Base Balance and Blood pH

Acid-Base Balance

• Metabolic processes produce: carbonic, sulfuric, phosphoric, and lactic acids.
• The bloodstream transports these to the lungs and kidneys, where they're expelled without affecting plasma pH because of: the blood's buffering capacity, respiratory mechanisms, and renal regulatory mechanisms.

Acids & Bases

• An acid is a proton (H+) donor
  • It donates hydrogen ions to a solution.
• A base is a proton (H+) acceptor
  • It accepts hydrogen ions from a solution.
• Strong acids dissociate quickly, releasing H+ ions.
  • Examples: HCl (hydrochloric acid), HNO3 (nitric acid), and H2SO4 (sulfuric acid).
• Weak acids partially dissociate in aqueous solutions.
  • An example is carbonic acid (H2CO3).
• Strong bases react quickly by removing H+ from a solution.
  • Hydroxyl groups (OH-) react with H+ to form H2O.
• Weak bases do not fully ionize in aqueous solutions.
  • An example is bicarbonate (HCO3-).
• Acid-base balance homeostasis mainly uses weak acids and bases.

Concept of Acidity

• The chemical activity of H+ ions defines a solution's acidity
  • H+ concentration determines blood acidity.
• The concentration of H+ ions is millions of times smaller than electrolytes like K+ and Na+ in body fluids, therefore :
  • K+ and Na+ concentrations are measured in milliequivalents per liter (mEq/L).
  • H+ is measured in nanoequivalents per liter (nEq/L).
• H+ ions are highly reactive and:
  • Body proteins contain many dissociable ions, like charged amino acids with side chains.
  • Amino acids gain or lose protons as H+ levels change, which affects protein structure and function.
• Maintaining constant [H+] in body fluids is crucial for enzyme and cellular function, for example :
  • Sodium-potassium ATPase pump activity decreases by half when pH drops by one unit.
  • Phosphofructokinase activity drops by 90% with a mere 0.1 unit pH decrease.

Concept of pH

• pH is the negative logarithm (base 10) of hydrogen ion concentration ([H+]).
  • This notation simplifies working with the wide range of [H+] values, in chemical systems.
• [H+] concentration determines the pH of body fluids
  • Homeostasis requires precise regulation of [H+], measured as pH (power of hydrogen).
• pH and [H+] change exponentially and not linearly.
  • Changing the pH by one unit means H+ concentration changes by a factor of 10.
• Normal extracellular fluid has a [H+] of 40 nEq/L, or 4 x 10-8 N.
• pH is 7.398 for blood and slightly lower (~7.2) intracellularly.
• Although [H+] is very low in the body it is very ‘powerful’
  • The H+ concentration, therefore, gets expressed as a logarithmic function called pH.
• pH is a numeric scale for expressing acidity or alkalinity of aqueous solutions.
  • Pure water (pH 7) is neutral and contains equal amounts of H+ (acid) and OH- (base).
• There is an inverse relationship between pH and [H+].
  • Higher [H+] means lower pH, leading to acidemia.
  • Lower [H+] means higher pH, leading to alkalemia.

Acidemia/Acidosis and Alkalemia/Alkalosis

• Acidemia is a decrease in blood pH below the normal limits.
• Acidosis refers to the pathophysiological processes that cause net acid accumulation.
• Alkalemia is an increase in blood pH above the normal limits.
• Alkalosis refers to the pathophysiological mechanisms resulting in alkali accumulation.
• Acidemia is defined as a pH less than 7.35 - 7.45 while alkalemia is pH greater than that

Concept of Buffering

• A buffer accepts or donates protons (H+), thus minimizing changes in pH.
• A buffer solution contains a weak acid and its conjugate base.
• Buffers reversibly bind H+ and prevent abrupt pH changes.
• Blood buffers are first line of defense that prevent major changes in blood pH

Law of Mass Action

• A chemical reaction's velocity is proportional to the product of reactant concentrations.
  • This law explains and predicts the behavior of solutions in a dynamic equilibrium.
• There are two opposing reactions for an acid, and the velocity of both reactions can be written as shown in image:
• If V1 = V2 it is a chemical equilibrium and the equilibrium constant is K.
• If H+ increases (acidemia), the system shifts left, consuming H+ and forming more HA.
  • Conversely, if pH increases, the reaction shifts right, releasing H+.

Henderson-Hasselbalch Equation and Isohydric Principle

• Body fluids use multiple buffering systems, where measuring every component is not necessary
  • Analyzing one system is enough to predict changes in others.
• The Henderson-Hasselbalch equation is a mathematical method in chemistry and biology to determine the pH of a buffer solution.
  • It describes the relationship between the pH and mixture of a weak acid and it's conjugated base
• The Isohydric principle states that multiple buffers in the same solution are always in equilibrium