Water Properties and Molecular Interactions

• Water's properties are fundamental to the behavior of biologic molecules.
• Water can form hydrogen bonds with itself (intramolecularly) and with other molecules it dissolves (intermolecularly).
• Water would be a gas without extensive intramolecular hydrogen bonds, differentiating it from other small molecules like CO2, CH4, NH3, O2, and N2.
• Hydrogen bonds are weak, reversible chemical bonds between molecules that can donate or accept partially charged hydrogen atoms.
• Intramolecular bonds in water create dynamically breaking and reforming tetrahedral structures.
• Water's hydrogen bonding affects the shape of surrounding biomolecules and can dissociate electrolytes into charged ions, which then associate to neutralize charges.

Hydrophobic and Hydrophilic Molecules

• Hydrophilic molecules are soluble due to their ability to form hydrogen bonds with water.
• Molecules forming many hydrogen bonds with water have higher solubility.
• Solubility decreases with increasing size due to water structure disruption, but large molecules maintain solubility through numerous hydrogen bonds.
• Hydrophobic molecules have low water solubility because they form few or no hydrogen bonds with water.
• Hydrophobic molecules aggregate to minimize water structure disruption.
• Water forcing hydrophobic molecules together is important for the three-dimensional structure of macromolecules and biological membranes.

Electrolytes

• Electrolytes dissociate into cations (positive charge) and anions (negative charge) when added to water, enabling electric current conduction.
• Strong electrolytes like HCl and NaCl dissociate completely in water.
• Weak electrolytes, typically organic acids like phosphoric and carbonic acids, do not fully dissociate and maintain equilibrium between undissociated (HA) and dissociated forms (A⁻).
• HA ⇄ H⁺ + A⁻
• Hydrogen ion concentration in a weak acid solution depends on the equilibrium constant (Keq) for the dissociation reaction
• HA ⇄ H⁺ + A⁻ ; Keq = [H⁺][A⁻] / [HA]
• Conjugate pairs serve as good buffers, resisting pH changes by re-establishing equilibrium when acid or base is added
• Increasing acidity shifts the equilibrium toward the undissociated form (HA), while decreasing acidity shifts it away from HA, restoring proton concentration.

Acids and Bases

• Acidic solutions contain more protons than produced by water ionization, while alkaline solutions contain fewer protons.
• Water ionization allows it to participate in weak acid equilibria.
• Strong electrolyte sodium acetate (CH3COONa) dissociates completely in water, producing acetate ions (CH3COO⁻).
• Acetate ions enter equilibrium with protons from water, reducing proton concentration and creating a slightly alkaline solution.
• Functional groups releasing free protons "act as" acids, while those accepting/binding them "act as" bases.
• Acids are proton donors, while bases are proton acceptors; acetate is the conjugate base of acetic acid.
• pH expresses proton concentration as a positive whole number rather than a negative exponent of 10: pH = -log[H⁺].
• pH units are exponents of 10, producing a logarithmic relationship with acidity, where increases in pH equate to reduced acidity.
• The pK value, a constant for an electrolyte, is the negative logarithm of the equilibrium constant, whereas pH changes with physiological conditions.
• Equilibrium constant for weak acid dissociation is termed Ka; the pK for an acid is defined as pKa.
• Acidic functional groups have pKa values below 7, while basic groups have pKa values above 7.

Henderson-Hasselbalch Equation

• When physiologic solutes modify a solution's pH, the equilibrium shifts the ratio of conjugate acids (HA) to conjugate bases (A-).
• The quantitative relationship between pH and the ratio of conjugate acid to base is described by the Henderson-Hasselbalch equation: pH = pKa + log([conjugate base]/[conjugate acid]) or pH = pKa + log([A-]/[HA]).
• For pH problems, set up the Henderson-Hasselbalch equation, fill in known values, and solve.
• Log (A⁻)/(HA) = logA⁻ – logHA.

Buffers and Titration Curves

• Buffers are conjugate pairs resisting pH changes, with their effect best shown by titration curves: plots of pH change upon strong base (e.g., NaOH) addition.
• pH is plotted (low to high), with an inflection point indicating buffering efficiency.
• The midpoint of the curve's inflection (pKa) reveals the smallest pH change with base addition, marking the best buffering range at pK ± 1 pH unit.

Carbonic Acid Conjugate Pair

• Carbonic acid (H2CO3) serves as a major acid-base buffer in blood, establishing equilibrium with volatile CO2 gas and bicarbonate ion (HCO3⁻).
• H2O + CO2 ⇄ H2CO3 ⇄ H⁺ + HCO3⁻
• Carbonic acid is often not included in the Henderson-Hasselbalch equation due to its rapid breakdown to bicarbonate or conversion to CO2 by carbonic anhydrase.
• Overall equilibrium between bicarbonate and CO2 relies on the rate of CO2 production in tissues and lung elimination, making lungs critical in regulating blood pH
• Impaired CO2 elimination from lung disease may lead to blood acidification (respiratory acidosis).
• Dissociation of a weak acid into a conjugate pair reaches midpoint with pH equaling pK for maximal buffering; Henderson-Hasselbalch equation links conjugate base-to-acid ratio to pH.
• Titration curves show an inflection point for each ionizable functional group; carbonic acid conjugate pair is in equilibrium with volatile CO2.

Acid-Base Properties of Amino Acids and Proteins

• Proteins gain their charge properties from amino acid side chains, with several capable of ionizing and acting as weak acids.
• The ionization, which depends on pK of the side chain's functional group, can produce a positive or negative charge.
• The Henderson-Hasselbalch equation describes ionization amount (dissociated-to-protonated ratio) for each group, ionizing independently with its pKa value.

Ionized Forms of Amino Acids

• A titration curve illustrates independent dissociation of a-amino and a-carboxyl groups.
• Protons remove from the carboxyl group first since it has the lowest pK (pKa = 2.3) before being removed from the amino group (pK = 9.9) as pH rises.
• Each pKa shows the midpoint of two equilibria, demonstrating buffering power in amino acids and proteins.
• At pH 7.0, amino acid side chains in proteins have these characteristic charges:
  • Positively charged: lysine, arginine.
  • Negatively charged: aspartate, glutamate.
  • Histidine: positively charged if pH drops below 6.0.
  • Cysteine: negatively charged if pH rises above 8.0.

Isoelectric Point

• Net charge on an amino acid or a protein equals the sum of charges on each amino acid side chain.
• The pH at which the molecule has a net zero (neutral) charge is its isoelectric pH (pI).
  • pH > pI: net negative charge.
  • pH < pI: net positive charge.
• Proteins do not migrate in an electrical field when the buffer's pH equals their isoelectric point, as there is no net charge to attract them.
• Side chains of amino acids (asp, glu, lys, arg, cys, his) act as weak acids at physiologic pH, giving proteins charge properties, and the isoelectric point represents the pH where the net charge sum is zero.