Chapter 2: The Chemical Foundation of Life – Weak Interactions in an Aqueous Environment

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Types and Strengths of Noncovalent Interactions

Overview of Noncovalent Interactions

Noncovalent interactions are fundamental to the structure and function of biomolecules. They are weaker than covalent bonds but essential for dynamic biological processes such as molecular recognition, protein folding, and enzyme catalysis.

Charge-charge interactions: Occur between fully charged particles (ions), often called salt bridges.
Dipole-dipole interactions: Occur between molecules with permanent dipoles.
Induced dipole interactions: Temporary dipoles induced by nearby charges or dipoles.
Hydrogen bonding: A special type of dipole-dipole interaction involving hydrogen atoms bonded to electronegative atoms.
van der Waals interactions: Weak, non-specific interactions due to transient dipoles.

Relative energies of chemical interactions

Relative Energies of Noncovalent Interactions

The strength of noncovalent interactions varies, influencing their biological roles.

Charge-charge: 13–17 kJ/mol
Hydrogen bond: 2–21 kJ/mol
van der Waals: 0.4–0.8 kJ/mol

Type of Interaction	Approximate Energy (kJ/mol)
Charge–charge	13–17
Hydrogen bond	2–21
van der Waals	0.4–0.8

Table of noncovalent interaction energies

Electrostatic Interactions and Dipoles

Permanent and Induced Dipoles

Dipoles arise from the separation of opposite electrical charges within a molecule. Permanent dipoles exist in molecules like carbon monoxide and water, while induced dipoles are temporary and result from nearby charges.

Permanent dipole: Molecules with uneven charge distribution (e.g., CO, H2O).
Induced dipole: Temporary shift in electron density due to external influence.

Permanent dipole in CO Permanent dipole in H2O

Types of Noncovalent Interactions

All noncovalent interactions are fundamentally electrostatic. Their strength depends on the nature of the interacting particles and their distance.

Type of Interaction	Model	Example	Dependence of Energy on Distance
Charge–charge	Ion pairs	NH3+ and COO–	1/r
Charge–dipole	Ion and dipole	NH3+ and H2O	1/r2
Dipole–dipole	Two dipoles	H2O and H2O	1/r3
Charge–induced dipole	Ion and induced dipole	NH3+ and benzene	1/r4
Dipole–induced dipole	Dipole and induced dipole	H2O and benzene	1/r5
Dispersion (van der Waals)	Transient dipoles	Benzene rings	1/r6

Table of noncovalent interaction types

Coulomb’s Law and Dielectric Constant

Coulomb’s Law in Biological Systems

The force between two charged particles is described by Coulomb’s Law. In biological systems, the medium (usually water) screens these interactions, reducing their strength.

Coulomb’s Law (vacuum):
Coulomb’s Law (in solution): , where is the dielectric constant.
Water’s dielectric constant: 80, making it an excellent solvent for ions.

Coulomb's Law in solution

van der Waals Interactions and Radii

van der Waals Forces

van der Waals interactions are weak, non-specific forces that arise from transient dipoles. They are significant when molecules are in close proximity.

van der Waals radii: The effective size of atoms or groups when interacting noncovalently.
Energy profile: Attraction increases as molecules approach, but repulsion dominates at very close distances.

van der Waals energy profile

Atom/Group	van der Waals Radius (Å)
H	1.2
O	1.4
N	1.5
C	1.7
S	1.8
P	1.9
–OH	1.4
–NH2	1.5
–CH2–	2.0
–CH3	2.0
Half-thickness of aromatic ring	1.7

Table of van der Waals radii

Hydrogen Bonding

Hydrogen Bond Characteristics

Hydrogen bonds are directional interactions between a hydrogen atom (bonded to an electronegative atom) and another electronegative atom. They are crucial for stabilizing protein and nucleic acid structures.

Donor: Atom to which hydrogen is covalently bonded.
Acceptor: Atom with lone pair electrons that interacts with hydrogen.
Bond length: Typically 2.8–3.1 Å.

Hydrogen bond donor and acceptor

Donor...Acceptor	Distance (Å)	Comment
O–H...O	2.8 ± 0.1	H bond formed in water
O–H...N	2.8 ± 0.1	Bonding of water to other molecules
N–H...O	2.9 ± 0.1	Important in protein/nucleic acid
N–H...N	3.1 ± 0.2	Relatively rare
O...H–N	3.7	Weaker

Table of hydrogen bond types

Properties of Water

Unique Properties of Water

Water is the medium of life due to its unique physical and chemical properties.

High heat capacity
High boiling point
High dielectric constant
Density greater in liquid than solid
Permanent dipole
Ability to form four hydrogen bonds

Iceberg showing density difference

Compound	Molecular Weight	Melting Point (°C)	Boiling Point (°C)	Heat of Vaporization (kJ/mol)
CH4	16.04	–182	–164	8.16
NH3	17.03	–78	–33	23.26
H2O	18.02	0	100	40.71
H2S	34.08	–86	–61	18.66

Table of water properties

Hydrophilic, Hydrophobic, and Amphipathic Molecules

Hydrophilic Molecules

Hydrophilic molecules interact favorably with water, often forming hydrogen bonds or ionic interactions. Water can compete with intramolecular hydrogen bonds, disrupting structures. Hydrophilic molecules in water

Hydrophobic Molecules

Hydrophobic molecules are nonpolar and do not interact favorably with water. Water forms clathrate structures around them, decreasing entropy and driving the hydrophobic effect, which stabilizes protein structures. Clathrate structure around hydrophobic molecule

Amphipathic Molecules

Amphipathic molecules contain both hydrophilic and hydrophobic regions. In aqueous solution, they can form monolayers, micelles, or bilayers, which are fundamental to membrane structure. Amphipathic lipid molecule Micelle and bilayer formation

Acids, Bases, and pH Calculations

Bronsted-Lowry Definitions

Acid: Proton donor
Base: Proton acceptor
Strong acid: Completely dissociates
Weak acid: Partially dissociates

Ionization of Water and pH Scale

Water ionizes to produce H+ and OH–. The ion product () is at 25°C.

pH calculation:
Neutral pH: 7 at 25°C

pH scale

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates pH, pKa, and the ratio of conjugate base to acid.

Equation:
Buffering range: pKa ± 1

Isoelectric Point (pI) and Net Charge Calculations

Isoelectric Point of Amino Acids

The isoelectric point (pI) is the pH at which the net charge of a molecule is zero. For amino acids like glycine, pI is calculated as the average of the two pKa values surrounding the neutral species.

Formula:

pI calculation for glycine Relative concentrations of glycine species

Biological Relevance of Ionic Properties

Ionic Properties of Amino Acid Side Chains

The ionic properties of amino acid side chains affect protein solubility, ligand binding, and purification processes.

Surface properties: Charged side chains are often solvent-exposed.
Protein-ligand binding: Electrostatic interactions are crucial for specificity.

Electrophoresis and Isoelectric Focusing

Agarose and Acrylamide Gel Electrophoresis

Agarose gel: Used for large molecules (DNA, RNA).
Acrylamide gel: Used for smaller molecules (proteins, short nucleotides).

Isoelectric Focusing

Proteins migrate in a pH gradient gel to the region corresponding to their pI, allowing separation based on net charge. Additional info: Isoelectric focusing is a powerful technique for protein purification and analysis, as it separates proteins based on their unique isoelectric points.