Acid-Base Balance and pH Regulation

Acid-Base Balance Importance

• Maintaining acid-base balance is important for body homeostasis, which is assisted by primary mechanisms

H+ Concentration

• H+ concentration is maintained in the body by the kidneys, lungs, and buffer systems

Relationship Between H+ and pH

• H+ ions are inversely associated with pH
• pH = -log [H+]
• pH is expressed as -log [H+] because H+ ion concentration is small
• If H+ is high, pH is low, so the solution is acidic
• If H+ is low, pH is high, so the solution is basic
• Base is synonymous with alkalinic
• Normal H+ concentration in blood is 40 nmol/L or 0.00004 mmol/L which is equal to a pH of 7.4

Acids and Bases Definition

• Acid molecules contain and can donate H+ ions in solutions
• Hydrochloric acid (HCl) forms hydrogen ions (H+) and chloride ions (Cl-) in water
• Carbonic acid (H2CO3) forms H+ and HCO3- (bicarbonate ions) in water
• A base is an ion or molecule that can accept H+
• Bicarbonate (HCO3-) can accept H+ to become H2CO3
• Proteins in the body often function as bases like haemoglobin
• Strong acids rapidly dissociate with H+ and release large amounts of H+ in solution
• Weak acids do not dissociate as rapidly
• Strong bases rapidly remove free floating H+ from solution
• Weak bases do not bind with H+ as rapidly

Normal pH Range

• Extracellular pH normally measures 7.4 with a range of 7.3 to 7.5
• Intracellular pH measures 6.0 to 7.4
• Intracellular pH is slightly lower than extracellular pH because the metabolism of cells produces acid
• Normal blood pH ranges between 7.35 and 7.45
• In alkalosis, there is excess removal of H+ from body fluids, so the pH increases
• In acidosis, there is excess addition of H+ in body fluids, so the pH decreases
• The bicarbonate buffer system is primarily responsible for maintaining blood pH
• A pH of less than 6.8 or greater than 8.0 results in death

Acid-Base Imbalances

• The body baseline is 20x more basic than acidic, so the body is slightly more alkalinic
• Acids are produced through ingesting food, and metabolism
• The body tries to restore pH back to normal through compensation when changes in normal pH occur
• If pH is not completely back in normal ranges, partial compensation is considered
• The body will continue towards complete compensation
• Homeostasis of pH is important for enzyme function and to speed up chemical reactions
• It is also important for electrolyte balance and concentration in the body and the function of certain hormones

Body Defenses to Changes in H+ Concentration

• Biological buffer systems defend against changes in H+ concentration

Chemical Buffer Systems

• Chemical acid-base buffer systems react within seconds to rapidly change pH
• Chemical buffers defend changes in pH of local microenvironments
• Buffers neutralize acids or bases
• Buffer systems consist of a weak acid and its conjugate base or a weak base and its conjugate acid
• Extracellular buffer system
• Bicarbonate buffer system is the most important
• Phosphate buffer system to a lesser degree
• Intracellular
  • Phosphate buffer system
  • Ammonia buffer system
  • Proteins

Bicarbonate Buffer System

• Bicarbonate buffer is the best
• It is associated with renal and respiratory systems
• It acts as both an extracellular and intracellular buffer
• Its weak acid = H2CO3 (carbonic acid)
• Its base = NaHCO3 (bicarbonate salt base)
• HCO3- : H2CO3 is maintained at a ratio of 20:1
• There is continuous motion between HCO3- and H2CO3 to keep a constant ratio
• When an acid is added: H+ + HCO3- -> H2CO3 -> CO2 + H2O
• The basic part of the buffer (bicarbonate) picks up the acid and makes carbonic acid
• The carbonic acid then creates carbon dioxide and water
• H+ is transferred to H2O so that the pH won't be affected
• When a base is added: NaOH + H2CO3 -> NaHCO3 + H2O
• Reaction occurs with the acid (carbonic acid) and it creates sodium bicarbonate and water
• The pH change won't be impacted

Carbonic Anhydrase

• Carbonic anhydrase is a family of metalloenzymes that catalyzes the reversible hydration of CO2 to HCO3- and H+
• This is the primary enzyme in the bicarbonate buffer system
• It is found in many tissues and involved in many systems
• It helps with CO2 transport for respiration and the reabsorption of HCO3- in the kidneys
• Helps maintain O2 binding in blood
• It is also used in the gastric system to produce gastric acid in the stomach; local response to maintain acid in stomach that is not actually part of acid base balance

Metabolic Acidosis in a Cell

• Ion transporters at the plasma membrane closely regulate pH inside cells
• Injecting hydrochloric acid (HCl) into the cytoplasm of a cell results in:
• Decreased pH in the cytoplasm due to H+ freely floating after disassociating from HCl
• Increased uptake of HCO3- into the cell
• Decreased movement of HCO3- outside of the cell
• Transport of H+ out of the cell
• Increase in overall pH due to fewer H+ floating around to decrease acidity

Metabolic Alkalosis Inside a Cell

• Injecting potassium hydroxide (KOH) base into the cytoplasm of a cell results in:
• Decreased uptake of HCO3- into the cell
• Increased movement of HCO3- outside of the cell
• A decrease in the exchange of H+
• Decreased pH because there are more H+ ions floating around to increase acidity

Phosphate Buffer System

• The phosphate buffer system is an intracellular and renal tubular fluid buffer
• It plays a minor role compared to HCO3- because there is not as much in the body
• But it plays a major role in the renal system occurring in proximal tubules, ascending loop of Henley, and early distal tubules
• Weak acid = H2PO4 (dihydrogen phosphate)
• Base = HPO4- (hydrogen phosphate)

Protein Buffers

• Proteins are known as 'Zwitter ions' because they can act as an acid in basic environments and as a base in acidic environments
• Intracellular fluid changes proportionately to extracellular fluid pH changes
• CO2 can rapidly diffuse through cell membranes and cause intracellular fluids to change with extracellular fluids
• Proteins within cells help to prevent changes
• Proteins include hemoglobi and plasma proteins
• During conversion of CO2 to HCO3- in red blood cells, H+ ions are buffered by hemoglobin (Hb)
• Acidic and basic amino acids, and carboxyl giving up H+, side chaisn with H+ are also present as buffers in plasma cell proteins in 27 amino acids

Chemical and Protein Buffers in Body Systems

• Biological buffers play a major role in the respiratory and renal systems
• The respiratory and renal system mostly maintain acid-base balance
• They do this through chemical and protein buffers

Respiratory System

• Main normal extracellular fluid pH is achieved by changing rate of breathing and maintaining constant PCO2
• CO2 is a byproduct of intracellular metabolism (cellular respiration) which affects extracellular pH
• Higher metabolism results in greater CO2 production
• Lower metabolism results in lower CO2 production
• The respiratory system senses changes in CO2 and adjusts ventilation accordingly
• The respiratory system H+ defense also relies on the bicarbonate buffer system

Bicarbonate Buffering in the Respiratory System

• CO2 readily combines with H2O in the presence of carbonic anhydrase to make H2CO3
• H2CO3 is a weak acid that readily dissociates in H+ and HCO3-
• Greater cellular respiration results in greater or lesserCO2 production and associated H+ changes
• The respiratory system can compensate by:
• Increasing ventilation to expire more CO2, which decreases [H+]
• Decreasing ventilation to retain more CO2 which increases [H+]

Renal Compensation

• Kidney is the most effective regulator of pH
• It takes longer than chemical buffers and the respiratory system
• Like the respiratory system, the kidneys also use bicarbonate buffering in the proximal convoluted tubule
• They also use phosphate buffer system in the tubular lumen
• There are three primary mechanisms for regulating H+ concentration: secretion of H+, reabsorption of filtered HCO3-, production of new HCO3-
• H+ secretion and HCO3- reabsorption occur in virtually all parts of the renal tubules except descending and ascending loops of Henle
• For each HCO3- reabsorbed, an H+ must be secreted
• 80-90% HCO3- is reabsorbed in the proximal tubule

Renal Compensation - Production of New HCO3-

• H+ ions combine with buffers in tubular fluids because enough isn't urinated in a single day to rid all the H+ ions free-standing
• The tubular fluids include ammonia and phosphate buffers
• Combining with other buffers allows for generation of 'new' HCO3- that is not being reabsorbed with every secretion of H+, This helps further buffer increases in H+ during acidosis

Respiratory Compensation

• Rapid response
• When [H+] increases and pH falls, increase respiration rate to remove more CO2
• When [H+] decreases and pH rises, decrease respiration rate to conserve more HCO3

Renal Compensation

• Slow response
• If pH falls when [H+] increases, more H+ is secreted, more HCO3- is reabsorbed
• If pH rises when [H+] decreases, secretion is decreased, and more HCO3- is secreted

Renal and Respiratory System Responsibilities

• The renal system is responsible for re-absorption of HCO3- and also excretion of non-volatile acids
• Re-absorption of HCO3- is necessary since it is the primary base in the body and needs to be reabsorbed after it is filtered out to keep the buffer system working properly
• Non-volatile acids are those which can't be exhaled like phosphoric acid, sulfuric acid, and ammonium
• The respiratory system is responsible for the removal of CO2 from the body
• CO2 is a volatile acid that can be exhaled or excreted as gas
• Chemical buffering plays a large role in both systems.

Acid-Base Imbalance

• The lungs and kidneys are used to compensate for dysfunction of the other to help maintain blood pH.
• An imbalance between the respiratory and renal systems can lead to alkalosis, acidosis, or become respiratory/metabolic.

Measuring Acid-Base Balance

• Arterial blood gases provide measure of: pH, PCO2, PO2
• Dissolved gases are used to calculated [HCO3-]

Anion Gap

• The concentrations in anions and cations in plasma must equal to maintain electrical neutrality
• Anions and cations are routinely measured including Anions (Cl- and HCO3-) and Cation (Na+)

Unmeasured Anions

• Albumin, phosphate, and sulfate
• Unmeasured Cations: calcium, magnesium, and potassium
• Anion gap = the difference between unmeasured anions and unmeasured cations
• It is used to diagnose different causes of metabolic acidosis which increases when unmeasured actions rise and measured actions fall

Pathologic Acid Excess

• In acidosis, excess production of a pathologic acid becomes an unmeasured anion buffered by HCO3-
• Lowering of HCO3- increases anion gaps and increases the unknown anion