Water, Hydrogen Bonding, Acids/Bases & Biological Buffers

Water as a Medium for Life

• Life originally evolved in water, partly because liquid water shields organisms from harmful UV radiation.
• Organisms are typically 70\%–90\% water by mass; in humans the range is 55\%–80\% of body weight.
• Virtually every biochemical reaction occurs in an aqueous environment; water therefore dictates reaction rates, equilibria, and product stability.
• Water is a critical determinant of macromolecular structure (proteins, nucleic acids, membranes) and thus their function.

Distribution of Water in the Body

• Intracellular fluid (ICF): fluid contained within cells.
• Extracellular fluid (ECF) has two sub-components:
  • Interstitial fluid (IF) located between cells.
  • Blood plasma circulating within vessels.
• Exchange of materials between plasma and cells always occurs through IF; capillary walls are permeable to water and small solutes.

Physiological Roles of Water

• Saliva formation and mucosal hydration (digestion and airway protection).
• Raw material for synthesis of hormones and neurotransmitters in the brain.
• Enables cell growth, reproduction, and survival by providing solvent and reactant.
• Major vehicle for excretion (urine) and thermoregulation (sweat/respiration).
• Lubricates joints, acts as a shock absorber for brain & spinal cord, transports O_2 and nutrients.
• Converts food to absorbable metabolites during digestion.

Thirst Mechanism

• Loss of body water ↓ blood volume → ↑ plasma osmolality → ↓ blood pressure.
• Detected by hypothalamic osmoreceptors and by renin–angiotensin system (↑ angiotensin II).
• Activates thirst center, causes sensation of dry mouth; water intake reverses osmolality changes.

Structure & Physical Chemistry of Water

Molecular Geometry and Dipole

• Obeys octet rule: O shares two e− pairs with two H, leaving two lone pairs.
• Geometry ≈ distorted tetrahedron; \angle\text{H–O–H} \approx 105^{\circ} (ideal sp^3 = 109.5^{\circ}).
• Oxygen is highly electronegative → large dipole moment; O atom carries partial \delta^-, H atoms \delta^+.
• Water therefore functions as both hydrogen-bond donor (via O–H) and acceptor (via O lone pair).

Hydrogen Bonding

• Defined as strong dipole–dipole interaction between a hydrogen bound to an electronegative atom (donor) and another electronegative atom (acceptor).
• Strength: 4–6\,\text{kJ mol}^{-1} (neutral) to 6–10\,\text{kJ mol}^{-1} (if one partner is charged) vs covalent O–H \approx 420\,\text{kJ mol}^{-1}.
• Directionality: maximal when donor, H, and acceptor are colinear.
• Each water molecule can form up to 4 H-bonds → explains high b.p., m.p., surface tension, cohesion & adhesion.
• Lifetime of an individual H-bond ≈ 1–20\,\text{ps}; network is highly dynamic ("flickering clusters").

Ice vs Liquid Water

• Hexagonal ice: each water fixed in lattice, 4 H-bonds exactly; lower entropy and density than liquid water (3.4 average bonds in liquid) → ice floats.

Solvent Properties of Water

Good Solvent For

• Polar & charged solutes: amino acids, peptides, carbohydrates, small alcohols, many salts.
• Mechanism: high dielectric constant screens ion–ion attractions, hydration by ion–dipole interactions lowers \Delta H and increases entropy \Delta S, giving negative \Delta G = \Delta H - T\Delta S.

Poor Solvent For (Hydrophobic)

• Non-polar gases (N2, O2), aromatic rings, aliphatic chains, typical waxes.
• Quantitative solubility table: e.g.
  • N_2 0.018\,g\,L^{-1} at 40^{\circ}C (nonpolar)
  • NH_3 900\,g\,L^{-1} at 10^{\circ}C (polar)

Biological Adaptations

• O_2 transported bound to Hb/Myoglobin (water-soluble proteins).
• CO2 converted to carbonic acid \mathrm{H2CO3} ⇌ bicarbonate \mathrm{HCO3^-} for aqueous transport.

Non-Covalent Interactions in Biology

• Ionic (Coulombic): charge–charge or charge–dipole.
• Hydrogen bonds: uncharged but polar.
• van der Waals: Universal attractive (London dispersion) + repulsive (steric) components; optimum at van der Waals contact distance \approx0.3–0.4\,\text{nm}.
• Hydrophobic effect: entropy-driven ordering of water around non-polar surfaces.

Hydrophobic Effect Detailed

1. Dispersed non-polar solute → water forms ordered "cage" (clathrate) → ↓ entropy.
2. Aggregation of non-polar parts (protein folding, lipid micelle formation) releases ordered water → ↑ entropy; net favorable.
3. Drives ligand binding: hydrophobic pockets in enzymes/receptors expel ordered water, stabilizing complex.

van der Waals Significance

• Individually weak but numerous; contribute to steric complementarity, DNA base stacking, binding of polarizable ligands.

Colligative & Non-Colligative Properties

• Colligative (depend only on particle number): b.p. elevation, f.p. depression, osmolarity.
• Non-colligative (depend on solute nature): viscosity, surface tension, taste, color.
• Cytoplasm is a concentrated solution → significant osmotic pressure.

Osmosis & Osmotic Pressure

• Water crosses semipermeable membrane from low [solute] to high [solute].
• At equilibrium, hydrostatic pressure balances osmotic driving force; pressure required defined as osmotic pressure \pi.
• Blood plasma osmolarity partly due to albumin & other proteins; changes affect cell volume.

Acids, Bases, and Ionization of Water

Definitions

• Acid = proton donor, Base = proton acceptor.
• Conjugate pairs: \text{HA} \rightleftharpoons \text{A}^- + \text{H}^+.

Auto-Ionization of Water

\mathrm{H_2O \rightleftharpoons H^+ + OH^-}

• Keq (at 25^{\circ}C) =1.8\times10^{-16}\,\text{M}.
• [\mathrm{H_2O}] = 55.5\,\text{M} (derived from density 1 g mL^{-1}).
• Ionic product Kw = K{eq}[\mathrm{H_2O}] = 1.0\times10^{-14}\,\text{M}^2.
• Therefore in pure water [\mathrm{H^+}] = [\mathrm{OH^-}] = 10^{-7}\,\text{M}.

Proton Hydration & Hopping

• Free \mathrm{H^+} does not exist; forms hydronium \mathrm{H_3O^+} and higher clusters.
• Proton hopping along hydrogen-bond network explains extraordinarily high mobility of \mathrm{H^+}.

pH Scale

\text{pH} = -\log[\mathrm{H^+}]

• Neutral: \text{pH}=7 when [\mathrm{H^+}] = [\mathrm{OH^-}].
• pOH analogous, with \text{pH} + \text{pOH} = 14 at 25^{\circ}C.
• pH can be <0 (e.g., [\mathrm{H^+}] = 6\,\text{M} \Rightarrow \text{pH} = -0.78).
• Examples: gastric juice 1–2; blood 7.35–7.45; 1 M NaOH ≈14.

Electrolytes

• Strong electrolytes dissociate fully (NaCl, HCl, NaOH).
• Weak electrolytes dissociate partially (acetic acid, NH_3).
• Non-electrolytes do not dissociate (sugars, alcohols, oils).

Weak Acids, pKₐ, and Titration Curves

Acid Dissociation Constant

K_a = \frac{[\mathrm{H^+}][\mathrm{A^-}]}{[\mathrm{HA}]}

• pKa = -\log Ka; stronger acid ⇒ larger Ka ⇒ smaller pKa.

Polyprotic Acids

• Donate >1 proton; each step has its own pK_a.
• Example phosphoric acid: pK{a1}=2.14, pK{a2}=6.86, pK_{a3}=12.4.
  • Predominant species at specific pH: at pH 2.14 \mathrm{H2PO4^-}, at pH 6.86 \mathrm{HPO4^{2-}}, at pH 10.5 mixture but mainly \mathrm{HPO4^{2-}}, at pH 12.4 \mathrm{PO_4^{3-}}.

Buffers

• Mixture of weak acid and its conjugate base resists pH change.
• Greatest buffering when \text{pH} = pK_a (acid & base 50:50).
• Effective buffer range: \text{pH} = pK_a \pm 1.

Henderson–Hasselbalch Equation

\text{pH} = pK_a + \log\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}

• Relates pH, pKₐ, and ratio of conjugate base to acid.

Example: Acetate Buffer

• \text{pK}_a = 4.76.
• Titration midpoint (pH 4.76) where [\mathrm{CH3COOH}] = [\mathrm{CH3COO^-}].

Comparative Titrations

• Acetic (pKₐ 4.76), phosphate (second pKₐ 6.86), ammonium (pKₐ 9.25) illustrate shifting buffering regions.

Biological Buffer Systems

• Intracellular: phosphate (millimolar); histidine side chains act near neutral pH.
• Blood plasma: bicarbonate/carbonic acid linked to respiration; pH maintained ≈7.4.
  • Equation: \mathrm{CO2 + H2O \leftrightarrow H2CO3 \leftrightarrow H^+ + HCO_3^-}.
  • Lung ventilation controls pCO_2, kidney excretion controls bicarbonate → integrated pH regulation.
• In vitro biochemical assays use zwitterionic sulfonic buffers: HEPES, PIPES, CHES etc.

Summary of Key Concepts & Connections

• Hydrogen bonding imparts water with unique physicochemical properties critical for life.
• Hydrophobic effect, van der Waals, ionic, and H-bond interactions collectively stabilize macromolecular 3-D structures and ligand binding.
• pH and buffering are governed by simple equilibrium principles (Henderson–Hasselbalch) yet underpin enzyme catalysis, metabolic regulation, and physiological homeostasis.
• Understanding water chemistry provides the foundation for analyzing biochemical pathways, drug design (hydrophobic pockets), and clinical scenarios (acid–base disorders, dehydration, osmotic imbalances).