Covalent and Noncovalent Bonds and Macromolecules

Covalent bonds

  • Covalent bonds are formed by sharing of electrons between atoms.
  • Key idea: sharing of valence electrons to achieve stable configurations.
  • Bond length: a specific distance between the nuclei of two atoms that is optimal for covalent bond formation. If atoms are too far apart, sharing is weak; if too close, nuclei repel each other.
  • Ball-and-stick representation: ball = nucleus, sticks = covalent bonds; used to illustrate 3D configurations.
  • Common atom bonding patterns:
    • Oxygen tends to form two covalent bonds (O often has 6 valence electrons and seeks two more).
    • Nitrogen tends to form three covalent bonds.
    • Carbon tends to form four covalent bonds.
  • Covalent bonds involve sharing electrons (valence electrons) between atoms.
  • Bond order and number of bonds:
    • A single covalent bond corresponds to two electrons being shared (one electron contributed by each atom).
    • A double covalent bond involves four electrons shared in total.
    • A triple covalent bond involves six electrons shared in total.
  • Bond strength vs stability:
    • Single bonds are the most stable (less reactive).
    • Double bonds are less stable/reactive than single bonds.
    • Triple bonds are the least stable among the three and require more energy to break.
    • Bond strength (energy required to break a bond) increases from single < double < triple.
  • Rotation around bonds:
    • Single bonds allow rotation around the bond axis.
    • Double and triple bonds restrict rotation, giving rigidity to the molecule.
  • Relevance to biomolecules:
    • Covalent bonds form the backbone of most biomolecules (e.g., carbon chains, rings).
    • Carbon, nitrogen, and oxygen frequently form covalent bonds in carbohydrates, proteins, nucleic acids, and lipids.
  • Summary key points:
    • Covalent bonds = sharing of electrons.
    • Bond length determines bond feasibility.
    • Bond order indicates number of shared electron pairs: single (2e−), double (4e−), triple (6e−).
    • Higher bond order = stronger bond but less rotational freedom (rigidity).

Electronegativity and polarity in covalent bonds

  • Electronegativity is the tendency of an atom to pull electrons toward itself in a covalent bond.
  • Not all atoms have the same electronegativity; electron distribution in a covalent bond can be unequal.
  • Polar covalent bond:
    • Unequal sharing of electrons creates partial charges:
    • Oxygen is more electronegative than hydrogen in H–O bonds, leading to b4 negative on O and b4 positive on H (represented as delta^− and delta^+).
  • Nonpolar covalent bond:
    • Atoms with similar electronegativity share electrons more equally; no significant charge separation.
  • Examples:
    • Carbon–Hydrogen (C–H): similar electronegativities -> nonpolar covalent bond.
    • Water (H$_2$O): O is more electronegative than H, causing polarity (O carries delta^−, H carries delta^+).
    • Oxygen–Oxygen (O=O) in O$_2$: same atom -> nonpolar covalent bond.
    • Carbon dioxide (CO$_2$): O more electronegative than C, giving C with delta^+ and O with delta^−.
  • Polar vs nonpolar is a continuum and affects solubility, interactions, and reactivity in biological systems.

Noncovalent bonds (four basic types)

  • Four basic types of noncovalent interactions:
    • Ionic bonds (electrostatic attraction between oppositely charged ions).
    • Hydrogen bonds (special kind of dipole–dipole interaction involving H attached to N, O, or F).
    • Van der Waals forces (nonspecific, weak attractions that arise when atoms come very close).
    • Hydrophobic interactions (not a true bond; tendency of nonpolar substances to aggregate in aqueous environments).

Ionic bonds

  • Formation: electron transfer from a donor atom to an acceptor atom, creating oppositely charged ions (cation and anion).
  • Example: Sodium chloride (NaCl).
    • Na loses its outer electron; Cl gains an electron.
  • Nature: electrostatic attraction holds the ions together.
  • Terminology: sometimes called electrostatic attraction.
  • Solubility: salts are highly soluble in water due to ion–water interactions.
  • Key terms: cation (positively charged), anion (negatively charged).

Hydrogen bonds

  • Definition: a relatively weak interaction that involves a hydrogen atom bonded to a highly electronegative atom (N, O, or F) forming an attraction with another electronegative atom elsewhere.
  • Key features:
    • One hydrogen atom is involved in the bond.
    • Commonly observed between water molecules and within biomolecules (DNA bases, protein secondary structures).
  • Examples in biology:
    • DNA base pairing: GC base pairs and AT base pairs stabilized by hydrogen bonds (GC forms three H-bonds; AT forms two H-bonds).
    • Protein structures: hydrogen bonds stabilize secondary structures like alpha helices and beta sheets.
  • Relative strength: weaker than covalent and ionic bonds but crucial for structure and function.

Van der Waals forces

  • Nature: nonspecific, weak attractions that arise from transient dipoles or induced dipoles in atoms/molecules.
  • Dependence: stronger for atoms with many electrons; stronger when atoms are closer together.
  • Relevance: contribute to condensation, packing of biomolecules, and interaction specificity when close in proximity.

Hydrophobic interactions

  • Not a true chemical bond; an emergent tendency for hydrophobic (water-fearing) molecules or regions to aggregate in aqueous environments.
  • Mechanism: in water, nonpolar groups disrupt hydrogen-bonding network of water; aggregation minimizes exposed surface area and stabilizes the system via entropy (hydrophobic effect).
  • Importance:
    • Drives folding of proteins by burying hydrophobic amino acids inside the core.
    • Critical for lipid organization and membrane structure.
  • Example: lipids (hydrophobic tails) aggregate in water; amphipathic molecules (e.g., fatty acids) have both hydrophobic and hydrophilic regions.

Hydrophilic vs hydrophobic molecules

  • Hydrophilic (water-loving): soluble in water; can form ionic bonds or hydrogen bonds with water.
    • Example: NaCl dissolves in water and dissociates into Na$^+$ and Cl$^-$; water stabilizes ions.
    • Example: urea (CO(NH$2$)$2$) can form hydrogen bonds with water despite lacking a formal charge.
  • Hydrophobic (water-fearing): insoluble in water; example: hydrocarbon chains.
  • Summary: solubility depends on ability to form bonds with water; hydrophilic molecules form ionic or hydrogen bonds with water, while hydrophobic molecules tend to aggregate via hydrophobic interactions.

Monomers and polymers (macromolecules)

  • Four major families of small organic molecules (monomers):
    • Sugars (carbohydrates)
    • Fatty acids (lipids)
    • Amino acids (proteins)
    • Nucleotides (nucleic acids)
  • Monomers join to form polymers/macromolecules via covalent bonds.
  • Polymers:
    • Carbohydrates: monosaccharides -> disaccharides -> oligosaccharides -> polysaccharides (e.g., glycogen).
    • Proteins: amino acids -> peptides -> proteins.
    • Nucleic acids: nucleotides -> nucleic acids (RNA, DNA).
    • Lipids: fatty acids + glycerol form larger lipids like triacylglycerols (storage fats).

Carbohydrates (sugars)

  • Monosaccharides: simplest sugars; cannot be hydrolyzed further.
  • Classification by carbon number:
    • Triose: 3 carbons
    • Tetrose: 4 carbons
    • Pentose: 5 carbons (e.g., ribose)
    • Hexose: 6 carbons (e.g., glucose)
  • Functional group classification:
    • Aldoses: carbonyl group is an aldehyde
    • Ketoses: carbonyl group is a ketone
  • Isomerism: same molecular formula, different arrangement (e.g., glucose, galactose, mannose).
    • Enzyme specificity: enzymes are stereospecific for exact isomers.
  • Oligosaccharides/polysaccharides:
    • Oligosaccharide: few sugars (<12)
    • Polysaccharide: many sugars (>12) (e.g., glycogen)
  • Glycosidic bond: covalent linkage between sugar residues via dehydration reaction (water removed).
    • Example disaccharides: maltose (glucose–glucose), lactose (galactose–glucose), sucrose (glucose–fructose).
  • Glycogen: branched glucose polymer; major energy storage molecule in liver and muscle; core concept: blue dots = glucose units; chains are branched.

Fatty acids and lipids

  • Fatty acids: long hydrocarbon chains with a terminal carboxylic acid group.
  • Length: 4 to ~30 carbons; classified as short, medium, long, very long chain.
  • Saturation:
    • Saturated fatty acids: no double bonds; straight chains (linear).
    • Unsaturated fatty acids: contain one or more C=C double bonds; cause kinks in chain and bending.
  • Amphipathic nature: fatty acids have a hydrophobic hydrocarbon tail and a hydrophilic carboxyl head; usually exist in stored form as triacylglycerol (TAG).
  • Triacylglycerol (triglyceride): storage form of fat; three fatty acids esterified to glycerol.
  • Bond linking fatty acids to glycerol: ester bond (formed by reaction between a carboxyl group of a fatty acid and a hydroxyl group of glycerol).
    • Ester bond: formed when R-COOH reacts with R'-OH and water is released.
  • Free fatty acids vs stored forms: in adipose tissue, fats are stored as TAGs.
  • Structural implications: presence or absence of double bonds changes shape and packing, influencing membrane fluidity and energy storage.

Amino acids and proteins

  • Amino acids: the building blocks of proteins.
  • Bond between amino acids: peptide bond (to be covered in detail in dedicated lectures).
  • Proteins fold and function due, in part, to noncovalent interactions that stabilize 3D structure (to be discussed later).

Nucleotides and nucleic acids

  • Nucleotides: monomers of nucleic acids (RNA and DNA).
  • Components of a nucleotide:
    • Nitrogenous base (adenine A, guanine G, cytosine C, thymine T, and uracil U in RNA)
    • Five-carbon sugar: ribose (RNA) or deoxyribose (DNA)
    • Phosphate group
  • Bond connecting nucleotides: phosphodiester bond (connects the 3' carbon of one sugar to the 5' carbon of the next sugar).
  • Nucleotides also serve as energy carriers and signaling molecules:
    • Energy carriers: ATP, GTP
    • Signaling molecules: cyclic AMP (cAMP), cyclic GMP (cGMP)

Bond types linking monomers to form macromolecules (summary)

  • Carbohydrates: glycosidic bond connects sugar residues (covalent).
  • Proteins: peptide bond links amino acids (covalent).
  • Lipids: ester bond links fatty acids to glycerol (covalent).
  • Nucleic acids: phosphodiester bond links nucleotides (covalent).

Structural basis vs functional interactions

  • Covalent bonds form the backbone of macromolecules, giving a stable 3D scaffold.
  • Noncovalent bonds and interactions (hydrogen bonds, ionic interactions, Van der Waals forces, hydrophobic interactions) are crucial for:
    • Determining the final three-dimensional structure of biomolecules.
    • Facilitating interactions between enzymes and substrates, receptors and ligands, and between nucleic acids and proteins.
    • Transient binding events in catalysis and signaling.
  • Enzymes and substrates interact through noncovalent contacts; the transient nature of these bonds allows binding and release during catalysis.

Quick reference: strength order (from strongest to weakest, as discussed)

  • Covalent bonds: greatest bond strength (highest energy to break) and form the backbone of macromolecules.
  • Ionic bonds: strong electrostatic interactions between oppositely charged ions.
  • Hydrogen bonds: weaker than ionic; crucial for structure and specificity.
  • Van der Waals forces: very weak, but collectively significant; strength scales with electron count and proximity.
  • Hydrophobic interactions: weak, non-bond-like attractions driving aggregation in water; important for folding and membrane structure.

Practical implications and connections

  • Understanding bond types helps explain:
    • Why biomolecules adopt specific shapes and how those shapes relate to function.
    • How mutations altering covalent backbones or noncovalent interaction patterns can affect stability and activity.
  • Relevance to medicine and biotechnology:
    • Drug design and design of enzyme inhibitors rely on mimicking or disrupting noncovalent interactions.
    • Protein misfolding diseases arise when noncovalent interactions fail to maintain proper structure.

Looking ahead

  • Next topics include enzymes and how they interact with substrates via noncovalent bonds, catalytic mechanisms, and energy changes.
  • We will explore structure–function relationships in proteins and nucleic acids in more depth.