Acids, Bases, and Acid-Base Reactions

• Early chemists identified acids as compounds with a sour taste, such as citric, acetic, and hydrochloric acid.
• Bases, or alkaline compounds, are substances that neutralize acids.
• Examples of alkaline solutions include wood ashes and glass cleaners.

Acids

• Acids are proton donors.
• Acidity is determined by the concentration of free hydrogen ions.
• Strong acids dissociate completely, releasing all H+ ions.
• Weak acids dissociate partially.

Bases

• Bases are proton acceptors.
• Strong bases dissociate rapidly and bind H+ ions almost immediately.
• Weak bases accept protons slowly.

Acid-Base Reactions

• Most acid-base reactions are reversible.
• A reversible reaction involves reactants (A and B) forming products (C and D), and products (C and D) reforming reactants (A and B).
• An irreversible reaction involves reactants (A and B) forming products (C and D), but products (C and D) do not reform reactants (A and B).
• The acid dissociation constant (Ka) indicates the extent to which an acid dissociates.

Ka vs pKa

• HCl has a Ka of 10^7, and acetic acid has a Ka of 1.74 x 10^-5.
• Acid strength can be indicated by its pKa value.
• pKa is calculated as -logKa.
• HCl has a pKa of -7, while acetic acid has a pKa of 4.74.

Acid Strength

• Very strong acids have a pKa less than 1.
• Moderately strong acids have a pKa between 1 and 3.
• Weak acids have a pKa between 3 and 5.
• Very weak acids have a pKa between 5 and 15.
• Extremely weak acids have a pKa greater than 15.
• A stronger acid has a larger Ka value.
• A stronger acid has a smaller pKa value.
• pH is calculated as -log[H+].
• pH indicates the concentration of protons in a solution.

Henderson-Hasselbalch Equation

• The Henderson-Hasselbalch equation: pKa = pH + log([HA]/[A-]).
• The Henderson-Hasselbalch equation can be used to determine whether a compound will be in its acidic form (with a proton) or basic form (without a proton) at a given pH.

Buffers

• Buffers resist changes in pH when acids or bases are added.
• Buffers are important for the proper functioning of cells and blood.
• Buffers maintain blood pH close to 7.4.
• Changes in blood pH affect oxygen uptake and cellular processes.

Buffers in Human Body

• Metabolic activity produces approximately 22,000 milliequivalents of acid per day.
• If all produced acid dissolved in water, pH would be less than 1.
• Blood pH typically ranges from 7.36 to 7.44.
• Intracellular pH is around 7.1 (6.9-7.4).
• Acid produced is buffered until excreted as CO2 through exhalation or as ions in urine.

Metabolic Buffers and Control Systems

• Chemical controls include bicarbonate, hemoglobin, phosphate, and protein buffer pairs.

• Physiological controls include the respiratory and renal control systems.

• 60% of a male body and 50% of a female body is water.

• 2/3 of body water is in intracellular fluid (ICF)

• 1/3 of body water is in extracellular fluid (ECF).

• Water input comes from food and ingested liquids.

• Water output occurs through water vapor, sweat, urine, and feces.

Bicarbonate Buffer System

• CO2 produced from the TCA cycle is the major source of metabolic acid in the body.

• The body generates 13 mol of CO2 per day.

• Carbonic acid has a pKa of 3.8 and is therefore almost completely dissociated.

• Though carbonic acid can be regenerated from CO2, theoretically, it cannot act as a buffer and generate bicarbonate.

• The addition of a base decreases [H] and [H2CO3], leading to more H2CO3 synthesized from dissolved CO2.

• Bicarbonate (HCO3-) and carbonic acid (H2CO3) are the base in this system.

• Measurement of carbonic acid (H2CO3) is not possible directly, its amount is estimated by assessing CO2's partial pressure because it has equilibrium with dissolved CO2

• CO2 can be dissolved.

• The rate of breathing and concentration is related.

Bicarbonate and Hemoglobin in Red Blood Cells

• The bicarbonate buffer system and hemoglobin in RBCs work together to buffer blood and transport CO2 to the lungs.
• Most CO2 diffuses into interstitial fluid and blood.
• Carbonic anhydrase in RBCs rapidly converts CO2 to carbonic acid.
• Increasing [H] from CO2 dissociation buffered by either hemoglobin or phosphate.
• Bicarbonate returned to the blood in exchange for Cl-.
• High bicarbonate levels in the plasma.
• Equilibrium shifts in the opposite direction approaching the lungs.
• CO2 released from RBC, causing more H2CO3 to dissociate.
• Hemoglobin loses some H+, increasing its affinity for oxygen.
• As pH decreases, the rate of breathing increases, and CO2 exhalation increases.
• Hypoventilation causes acidosis.
• Hyperventilation causes alkalosis.

Hemoglobin Buffer

• Hemoglobin buffers H+ from metabolically produced CO2 in the plasma.
• Hemoglobin has a greater affinity for H+ as it releases O2.
• H+ generated at the tissue level from the dissociation of H2CO3.
• Bound H+ to Hb does not contribute to blood acidity.
• As H+Hb picks up O2 from the lungs, Hb releases H+ and picks up O2.
• Liberated H+ combines with HCO3-.
• Venous blood is only slightly more acidic than arterial blood due to Hb's buffering capacity.

Other Acids

• Other acids include acetoacetic acid, β-hydroxybutyric acid, and lactic acid.
• These acids have a pKa of approximately 5 and are completely dissociated in blood and cellular fluid.
• Metabolic anions are transported out of the cell together with H+.

Urinary Hydrogen, Ammonium, and Phosphate Ions

• The nonvolatile acids produced in the body must be excreted in urine.

Phosphate Buffers

• A buffer of both ICF and urine.
• Phosphate anions and proteins are the major buffer systems of intracellular fluid.
• H2PO4 generates H+ ions.
• They play a major role in RBCs and all other cells.
• ICF has a high concentration of proteins that behave as H+ ion acceptors.

Urine

• Urine pH is between 5.5 and 7.
• Phosphate is excreted in urine.
• Ammonia (NH3) is a base that combines with protons to generate ammonium (NH4+).
• Ammonia is produced from the digestion of amino acids and is kept at low concentrations in the blood.
• Cells in the kidney generate NH4+ and excrete it into the urine in proportion to the acidity of the blood.

Protein Buffer System

• Proteins are very large, complex molecules.
• Proteins surrounded by negative charges on the outside and positive charges in the crevices of the molecule.
• H+ ions are attracted to, and held from, chemical interaction by the negative charges.
• OH- ions, which are the basis of alkalosis, are attracted by the positive charges in the crevices of the protein.

Chemical Control

• Bicarbonate, hemoglobin, phosphate, and protein buffer pairs are responsible for chemical control of pH.

Respiratory Control System

• The respiratory control system regulates pH by adjusting the rate and depth of breathing.
• Hypoventilation leads to acidosis.
• Hyperventilation leads to alkalosis.

Renal System

• The renal system regulates pH by excreting acids and bases in the urine.

• Different buffer systems assume dominant roles related to the body part.

• Extracellular Fluid is mainly buffered by Bicarbonate, but also Intracellular proteins and Phosphate.

• Blood is mainly buffered by Bicarbonate and Hemoglobin, but also Plasma proteins and Phosphate.

• Intracellular Fluid is mainly buffered by Proteins and Phosphate.

• The main buffers in urine are Ammonia and Phosphate.

• In acid-base imbalance, respiratory alkalosis is caused by either hyperventilation or decreased CO2.

• Blood pH is termed acidosis if it is < 7.35.

• The most effective buffer system at physiological pH is the Bicarbonate buffer.

• The greatest buffering capacity at physiological pH would be by a protein rich in Histidine.