Biomolecular Interactions and Water as the Medium for Life

Intermolecular Interactions in Biomolecules

Primary theme: biomolecules are stabilized and shaped by a set of weak, dynamic interactions. These interactions are powerful in aggregates and structures (proteins, membranes) but individually are weak compared with covalent bonds.
The four main weak interactions discussed:
- Hydrophobic interactions (HP)
- Hydrogen bonds (HB)
- Ionic (electrostatic) interactions
- Van der Waals (VDW) interactions (London dispersion and induced dipole interactions)
Relative strengths (approximate, per mole):
- Hydrophobic interactions: $E_{HP} \gtrsim 20\ \text{kJ mol}^{-1}$ (strongest among the four listed)
- Ionic/electrostatic interactions: $E_{ionic} \gtrsim 20\ \text{kJ mol}^{-1}$
- Hydrogen bonds: $E_{HB} \approx 12-20\ \text{kJ mol}^{-1}$
- Van der Waals: $E_{VDW} \approx 4\ \text{kJ mol}^{-1}$ (roughly a few kJ/mol, weak)
Covalent bonds are much stronger: typically > 300\ \text{kJ mol}^{-1} (single bonds ~ 350–400 kJ/mol, up to ~900 kJ/mol for multiple bonds)
Overall hierarchy (strength) among the four weak interactions: HP > Ionic > HB > VDW
The weak interactions are essential for structure and function but are also dynamic and reversible; they constantly break and re-form on fast timescales.
Dynamics and timescales:
- The interactions are dynamic and constantly rearranging, especially within proteins and in aqueous environments.
- Picosecond-scale fluctuations: typical time scale mentioned for internal rearrangements is around $\sim 10^{-12}\ \text{s}$ (picoseconds).
Functional significance:
- Structural stabilization: the network of weak interactions holds secondary, tertiary, and quaternary structures together.
- Folding and recognition: hydrophobic collapse drives core formation; long-range interactions pull distant regions into a functional conformation.
- Dynamics enable conformational changes necessary for activity; disrupting a single weak interaction can disrupt function.
Role of hydrophobic interactions in biological membranes and early life:
- Hydrophobic amino-acid stretches tend to cluster to shield nonpolar parts from water, driving proper folding.
- In membranes, hydrophobic tails face the lipid interior, while polar heads face water, creating a lipid bilayer.
- Complex lipids have polar (hydrophilic) heads and nonpolar (hydrophobic) tails; HP interactions help keep the bilayer intact.
- The lipid bilayer is presented as fundamental to the evolution of cellular life and the formation of a compartmentalized membrane (foundation of the cell).
Covalent interactions mentioned for proteins:
- A covalent linkage (often referred to as a disulfide “sulfa linkage” in the lecture) can stabilize structure and is discussed as something to be covered later.

Hydrophobic Interactions and Their Roles

Hydrophobic interactions arise when nonpolar (hydrophobic) molecular segments aggregate to minimize contact with water.
In a polar aqueous environment, hydrophobic segments tend to come together, increasing the entropy of water by freeing structured water around exposed nonpolar surfaces.
Consequences:
- Formation of a hydrophobic core in folded proteins.
- Stabilization of lipid bilayers, where nonpolar lipid tails are shielded from water.
- Driving force behind membrane assembly and the separation of cellular compartments.
Hydrophobic interactions are a key part of how complex lipids maintain a bilayer by burying nonpolar interiors and exposing polar heads to water.

Hydrogen Bonds (HB)

Hydrogen bonds form between a hydrogen atom covalently bound to an electronegative atom (e.g., N, O) and another electronegative atom with lone pairs.
In water, each molecule can form hydrogen bonds with several neighbors, creating an extensive H-bond network.
Energy range and significance:
- Hydrogen bonds typically contribute about $E_{HB} \approx 12-20\ \text{kJ mol}^{-1}$ depending on environment and participating atoms.
Structural role in proteins:
- Backbone amide hydrogen bonds stabilize secondary structures such as alpha helices and beta strands.
In water:
- Water forms a dynamic, extensive hydrogen-bonded network; a water molecule usually participates in multiple hydrogen bonds with surrounding molecules.
- The hydrogen-bond lifetime is short; a typical half-life is about $t_{1/2} \approx 9.5\ \text{ps}$ for hydrogen-bond exchange.

Van der Waals (VDW) Interactions

Van der Waals interactions include dispersion forces (London forces) and induced dipole-induced dipole interactions.
They can operate even within a single group of atoms (e.g., alkyl groups) and contribute to fine-tuning of molecular packing.
Role in proteins:
- Contribute to stabilization of folded conformations alongside HB and HP interactions.
Relative weakness:
- Generally weaker than HB and ionic interactions, but collectively important for close-contact packing and specificity.

Water as the Medium for Life and Its Unique Properties

Water is a polar molecule with a bent geometry: the H–O–H angle is about $104.3^{\circ}$ (often cited as ~109.5° for ideal tetrahedral; here 104.3° reflects lone-pair repulsion).
Polar nature and hydrogen-bonding network:
- Water forms a hydrogen-bonded network that grants unique solvent properties important for biomolecules.
Dielectric constant:
- Water has a high dielectric constant, which reduces electrostatic attraction between ions in solution and stabilizes dissolved ions via hydration shells.
- Example: NaCl in water dissociates into hydrated Na+ and Cl− ions with hydration shells that prevent ions from re-associating readily.
Hydration shells:
- Cations are surrounded by water molecules (hydration shells) that partially neutralize the charge and hinder close approach of ions.
- The hydration effect stabilizes ions and contributes to the high dielectric environment.
Solvent properties and solubility:
- Because of its polarity, water dissolves many biomolecules (proteins, nucleic acids, sugars) through hydrogen bonding with polar groups.
- Sugars (e.g., glucose) with many hydroxyl groups form extensive hydrogen-bond networks with water, contributing to high solubility.
- Lipids contain both polar and nonpolar regions; their solubility in water is governed by amphipathic character and hydrophobic segregation.
Amphiprotic/antiprotic properties of water:
- Water is amphiprotic: it can act as an acid or a base, donating or accepting a proton:
- In acid-base terms: $\mathrm{H2O} \rightleftharpoons \mathrm{H3O^+} + \mathrm{OH^-}$
- Water can act as an acid (donating a proton to become ( \mathrm{OH^-} )) or as a base (accepting a proton to become ( \mathrm{H_3O^+} )).
Autoionization constant and pH concepts:
- The autoionization constant for water is denoted as $Kw = [\mathrm{H^+}][\mathrm{OH^-}]$ , commonly approximated as $Kw \approx 1.0 \times 10^{-14}$ at 25°C.
- In pure water at 25°C: $[\mathrm{H^+}] = [\mathrm{OH^-}] = 10^{-7} \text{ M}$ , so $\mathrm{pH} = \mathrm{pOH} = 7.$
pH and pOH relation:
- The fundamental relation: $\mathrm{pH} + \mathrm{pOH} = 14.$
Buffering and the limitations of pure water as a buffer:
- Water is a poor buffer because its pH changes abruptly with small additions of acid or base due to the simple equilibrium and the fixed concentration of water.
- Henderson–Hasselbalch framework helps design buffers to resist pH changes:
- For a weak acid HA and its conjugate base A−, the Henderson–Hasselbalch equation is:
- $\mathrm{pH} = \mathrm{p}K_a + \log \left(\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}\right)$
Weak acids and buffer formation:
- For a weak acid HA with conjugate base A−, the dissociation constant is:
- $Ka = \frac{[\mathrm{H3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$
- $\mathrm{p}Ka = -\log Ka$
- Rearranging the above, you can derive the Henderson–Hasselbalch equation and understand buffer behavior under acid/base addition.
Example of buffer behavior and capacity:
- If the ratio [A−]/[HA] equals 10, the pH equals $\mathrm{p}Ka + 1$ ; if the ratio is 1/10, the pH equals $\mathrm{p}Ka - 1$ . This shows how buffers resist pH changes within a range around the $\mathrm{p}K_a$.
- Buffer capacity describes how much acid or base a buffer can neutralize with only a small change in pH. It depends on the absolute amounts of HA and A− and their ratio.

Acid–Base Chemistry of Water and Its Relevance to Biomolecules

Water as amphiprotic: can donate or accept a proton, enabling it to act both as an acid and as a base.
pH changes and biological macromolecules:
- Small pH shifts can disrupt ionic interactions within proteins (e.g., protonation/deprotonation of amino and carboxyl groups affecting ionic bonds and salt bridges).
- Such disruptions can alter protein folding, stability, and function.
Buffer systems in biology are essential for maintaining pH within narrow windows suitable for enzyme activity and macromolecular stability.

Implications for Protein Structure and Function (Integrated View)

Structural integrity relies on a balance of interactions:
- Hydrophobic core formation and stabilization by HP interactions.
- Backbone hydrogen bonding supporting secondary structures.
- Electrostatic and ionic interactions (salt bridges) helping stabilize tertiary structure and folding patterns.
- Van der Waals interactions fine-tune packing and specificity.
Chemical alterations (e.g., changes in pH) can alter the balance by changing protonation states of ionizable groups (e.g., NH3+ vs NH2, COOH vs COO−), breaking ionic interactions and potentially altering or destroying native structure.
In a broader sense, water’s properties (polarity, high dielectric constant, hydrogen-bond network, and buffering capacity) provide the environment in which biomolecules fold, interact, and function.

Key Equations and Concepts to Remember

Hydrogen bond energy range and comparison:
- $E_{HB} \approx 12-20\ \text{kJ mol}^{-1}$
Hydrophobic interaction energy scale (relative):
- $E_{HP} \gtrsim 20\ \text{kJ mol}^{-1}$
Ionic interactions: $E_{ionic} \gtrsim 20\ \text{kJ mol}^{-1}$
Van der Waals interactions: $E_{VDW} \approx 4\ \text{kJ mol}^{-1}$
Relative strength order: HP > Ionic > HB > VDW
Core relationships:
- Water autoprotolysis: $K_w = [\mathrm{H^+}][\mathrm{OH^-}] \approx 1.0\times 10^{-14}$ at 25°C
- In water: $[\mathrm{H^+}] = [\mathrm{OH^-}] = 10^{-7}\ \text{M}$ and $\mathrm{pH} + \mathrm{pOH} = 14$
Acid–base definitions:
- $Ka = \frac{[\mathrm{H3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$
- $\mathrm{p}Ka = -\log Ka$
- Henderson–Hasselbalch: $\mathrm{pH} = \mathrm{p}K_a + \log \left(\frac{[\mathrm{A^-}]}{[\mathrm{HA}]}\right)$
Water as solvent and hydration shells:
- Hydration shells: cations are surrounded by water molecules, reducing their mobility and stabilizing charged species in solution.
- Water’s polarity enables extensive hydrogen bonding with polar biomolecules (nucleic acids, proteins, carbohydrates).

Quick Takeaways for Exam Preparation

The stability and function of biomolecules arise from a delicate balance of weak, dynamic interactions (HP, HB, ionic, VDW), with HP often providing the driving force for hydrophobic core formation and membrane assembly.
Water’s properties (polarity, high dielectric constant, hydrogen-bond network, amphiprotic behavior) are central to biomolecule solubility, structure, dynamics, and function.
pH and buffering are essential for maintaining physiological conditions; Henderson–Hasselbalch provides a practical framework to predict buffer behavior and capacity in the presence of weak acids/bases.
Small changes in pH can disrupt ionic interactions in proteins, highlighting why buffers and regulated cellular pH are critical for enzymatic activity and structural stability.