pH Balance and Buffer Systems

Lecture Learning Outcomes

• Buffer systems maintain pH stability in biological and chemical systems
• The Henderson-Hasselbalch equation interprets acid/base equilibria
• The Henderson-Hasselbalch equation determines the ionisation state of weak acids and bases given the pKa and pH
• Buffer capacity is influenced by factors particularly in biological systems
• The bicarbonate buffer system maintains physiological pH balance.
• Acidosis and alkalosis intricacies beyond the scope of the session

Maintaining pH

• pH measures hydrogen ion concentration (pH = -log[H+]) on a log scale
• Neutral pH (pH=7) equates to 10-7 M of H+
• Small pH changes represent large [H+] changes due to the log scale
• Body compartments maintain 'normal' pH within tight limits
• 'Normal' pH varies across body compartments like arterial blood (7.41), venous blood (7.36), intracellular fluid (7.35), and urine (~6-8)
• Changes in intracellular fluid pH affect extracellular pH and vice versa
• Problems occur rapidly when pH moves outside the normal range
• Proteins and enzymes denature, which leads to metabolism malfunction
• Acidosis or alkalosis are life-threatening if untreated with a blood pH below 7.35
• The CNS malfunctions and coma may result in blood pH falling below 7.35

Buffers Explained

• A buffer is a mix of weak acid/base and its conjugate form
• Weak acid buffers are aqueous solutions of a weak acid (HA, e.g., acetic acid) and its conjugate base salt (e.g., NaA, sodium acetate)
• Equilibria have certain concentrations of HA and A-
• Changes in pH drive the system to counteract the change based on Le Chatelier's principle
• The Henderson-Hasselbalch equation restates the acid dissociation constant, Ka, in terms of pH

Buffer Action

• Weak acids in the body are generally organic, such as carboxylic acids or inorganic, such as phosphoric acid
• Buffers resist pH changes upon addition of H+ or OH-
• When pH equals pKa, [HA] equals [A-]
• Buffers function best within +/-1 pH unit from the pKa, with about 10-90% dissociation

Buffer Capacity

• Buffering ability relies on concentration
• When the weak acid (HA) and conjugate base (A-) are used up the ability to resist pH change exhausts
• Greater availability of A- and HA molecules reduces the impact of strong acid (H+) or strong base (OH-) addition on pH
• Body buffers determine ability to cope with acidity changes
• Kidneys supply bicarbonate (HCO3-) to the blood
• Bicarbonate mops up H+
• Metabolic acidosis can result from improperly functioning kidneys.

Need for Buffers

• Most biological processes are pH-dependent, with enzymes functioning in a narrow pH range
• Small pH changes drastically affect reaction rates and system integrity
• Normal metabolism constantly generates CO2 and H+, disturbing pH
• Buffers maintain a suitable physiological pH

Body Buffers

• Phosphate buffer is intracellular with H2PO4- ↔ HPO4- + H+, pKa = 6.86
• Bicarbonate buffer system is in extracellular fluid (ECF), red blood cells (RBC), and blood; it has reversible equilibria including dissolved CO2
• Proteins are found throughout the body, such as albumin in blood and haemoglobin

Body buffering

• The body uses chemical buffering to minimize changes in [H+], but it does not remove it
• Body buffering involves food intake, digestion, and absorption with normal net acid input of ~1mM/Kg body weight/day
• The body employs cell metabolism, respiratory response, and renal response with chemical buffering of sulphate and phosphate
• Chemical buffering contains Chloride while binding to bases and excreting urine.

Phosphate Buffer System

• The main intracellular buffer alongside proteins with about 6X higher concentrations in ICF compared to ECF
• At physiological pH, the dissociation of the 2nd H+ is relevant at a pKa of H2PO4- of 6.86
• Buffers work best between +/- 1 from the pKa and act as effective buffers at a pH ~ 7.4

Bicarbonate System

• Key blood buffer is maintained at nearly constant pH (~pH 7.4; venous ~7.36)
• Also a key intracellular buffer in RBCs.
• Aids cell volume and fluid secretion maintainence and eliminates CO2 (major waste product).
• Consists of kidneys which excrete or retain H+ in the form of ammonia (NH4+) and lungs which exhale/inhale volatile acid (CO2)
• [CO2] can be rapidly adjusted by changes in rate of breathing with a large reservoir of CO2
• Bicarbonate (HCO3-) transport proteins provide pH inside and outside of RBCs
• Diseases like cystic fibrosis and heart disease may result due to disruption of this transport system

Bicarbonate Equilibria

• A drop in blood pH results in lactic acid production or exogenous salicylates
• Equilibria shifts towards production of CO2(g) increasing its exhalation
• Elevated NH3 levels cause reduction in cellular H+ resulting in blood pH rising
• Causes plasma [H+] to decrease,
• H2CO3 dissociates causing more CO2(g) in lungs to dissolve.

H-H Bicarbonate Equation

• Normal levels of bicarb [HCO3-] are ~24mM with a carbonic acid [H2CO3] of ~1.2mM
• [H2CO3] is expressed as pCO2 with dCO2 = 0.03 x pCO2

Respiratory vs. Metabolic Control

• [H2CO3] is normally ~1mEq/L

• CO2 released during respiration limits supply of H2CO3 with volatile acid (CO2) exhaled from lungs

• [H2CO3] is under respiratory control

• Respiratory function impairments cause acid/base imbalance

• [HCO3-] is normally ~25mEq/L and closely regulated by kidneys

• Kidneys are ultimate acid/base regulators and are acting slowly to compensate for disturbances arising from diet, metabolic, or disease

• Body buffers temporarily tie up acids while the kidneys permanently excrete them

• Metabolic (fixed) acids arise from incomplete metabolism (lactate, ketones) and proteins (sulfate)

Acid/Base Balance

• Concentration of buffering species determines the body’s buffering capacity
• Rate of acid/base production exceeding buffering species’ ability overwhelms the system
• In the bicarbonate buffer system, this is called the alkaline reserve (available HCO3-)

Acid Excretion

• Excess acid is transferred to NH3 (a weak base) in the kidney's proximal convoluted tubule (PCT) forming ammonium, NH4+
• Ammonium is very weak acid which stays protonated and is excreted in urine
• Simultaneously HCO3- is produced in the kidneys and transferred into the blood
• Kidneys replenish the HCO3- buffering capacity
• Respiratory/metabolic acid/base imbalances result from kidney or lung malfunction