Molecular Structure & Biological Macromolecules

The properties of biological molecules -- how they interact, fold, and function -- are ultimately determined by chemistry: the types of bonds that hold atoms together, the geometry those bonds impose, and the non-covalent forces that drive assembly of complex structures. This lesson examines covalent and non-covalent bonding, then applies these principles to understand the four major classes of biological macromolecules: lipids, proteins, carbohydrates (briefly), and nucleic acids.

Covalent Bonds arise from the sharing of electron pairs between atoms. Bond strength is characterised by bond dissociation energy (BDE): C-C ~= 346 kJ/mol; C=C ~= 611 kJ/mol; C?C ~= 837 kJ/mol; O-H ~= 459 kJ/mol. Electronegativity differences between atoms create polar covalent bonds with partial charges (delta+ and delta-). The C-O and O-H bonds in organic molecules are polar, making alcohols, carboxylic acids, and amino groups capable of hydrogen bonding. Bond geometry is predicted by VSEPR theory: carbon with four single bonds adopts tetrahedral geometry (109.5degrees); sp2 carbon in double bonds is trigonal planar (120degrees); the peptide bond has ~40% double-bond character due to resonance, giving it planar rigidity.

Non-Covalent Interactions are individually weak but collectively powerful -- the forces driving protein folding, enzyme-substrate recognition, membrane assembly, and DNA double helix formation.

Hydrogen bonds form between a hydrogen atom covalently bonded to an electronegative atom (N, O, F -- the donor) and a lone pair on another electronegative atom (the acceptor). Strength: ~20 kJ/mol. DNA base pairing (A-T: 2 H-bonds; G-C: 3 H-bonds), alpha-helix and beta-sheet formation in proteins, and water's high boiling point all depend on hydrogen bonds.

Van der Waals forces (London dispersion forces) arise from transient fluctuating dipoles in any two closely apposed atoms. Each interaction is very weak (~0.4-4 kJ/mol) but they become significant when large surfaces are complementary -- explaining why hydrophobic protein interiors pack tightly and why lipid bilayer leaflets associate.

Ionic (electrostatic) interactions between oppositely charged groups (e.g., lysine -NH3+ and aspartate -COO-) in proteins stabilise tertiary structure and enzyme active-site interactions. These are strongly influenced by pH and ionic strength.

The hydrophobic effect is the dominant driving force for protein folding and membrane assembly. Non-polar groups cannot form hydrogen bonds with water and instead cause ordering of the water shell (an entropically unfavourable state). Clustering of hydrophobic residues in a protein's core or lipid tails in a bilayer releases ordered water molecules into the bulk, increasing entropy -- a thermodynamically favourable process (deltaG < 0 because of the TdeltaS term).

Lipids are a chemically diverse class of hydrophobic or amphipathic molecules. Phospholipids spontaneously self-assemble into bilayers -- the thermodynamic basis of all biological membranes. Cholesterol intercalates between phospholipids, moderating membrane fluidity. Triglycerides store energy.

Proteins at the chemical level are polymers of amino acids joined by peptide bonds. The backbone conformation is constrained by the planarity of the peptide bond (resonance) and steric limits described by the Ramachandran plot. alpha-helices are stabilised by intra-chain backbone hydrogen bonds; beta-sheets by inter-strand backbone hydrogen bonds.

Nucleic Acids. DNA and RNA are polynucleotide chains. Base stacking and Watson-Crick hydrogen bonds (A:T -- two H-bonds; G:C -- three H-bonds in DNA; A:U in RNA) stabilise the double helix. The G:C pair's extra hydrogen bond explains why GC-rich DNA regions require more energy to denature. The 2'-OH group of ribose in RNA makes it susceptible to hydrolysis, explaining why RNA is less stable than DNA.