Chemical Foundations and Molecular Interactions

Chemical Foundation

• Covalent bonds form when atoms share electrons, creating stable bonds.
• Nonpolar covalent bonds share electrons equally, like in hydrocarbons.
• Polar covalent bonds share electrons unequally, creating partial positive and negative charges (dipoles), as seen in water.
• Non-covalent interactions are weaker than covalent bonds, essential for transient biological interactions.
• Ionic bonds occur when oppositely charged ions attract, involving complete electron transfer (e.g., Na+ and Cl-).
• Hydrogen bonds involve a partially positive hydrogen atom and an unpaired electron pair (often with O or N) forming weak bonds.
• Van der Waals forces are weak, non-specific attractions between close atoms, important for molecular packing and shape.
• Hydrophobic interactions occur when nonpolar molecules cluster together in water to minimize energy.

Molecular Complementarity

• Molecules interact based on complementary shapes, charges, and other properties.
• This interaction facilitates dynamic biological processes through specificity and affinity.
• Specificity is the precision of fit between interacting molecules.
• Affinity describes the strength of binding, measured by the dissociation constant (Kd).
• Examples include receptor-ligand binding and antibody-antigen interactions.
• Proteins can have multiple binding sites for various ligands.

Chemical Building Blocks

• Macromolecules are constructed from monomers via condensation reactions (removal of water).
• Proteins are composed of amino acids linked by peptide bonds.
• Nucleic acids are made of nucleotides linked by phosphodiester bonds.
• Carbohydrates are built from monosaccharides joined by glycosidic bonds.
• Lipids are non-polymeric, often composed of hydrophobic tails (e.g., fatty acids) and hydrophilic heads.
• Condensation reactions combine monomers by removing water.
• Hydrolysis breaks down macromolecules by adding water.

Chemical Reactions and Equilibrium

• Chemical equilibrium occurs when the rates of forward and reverse reactions are equal.
• Equilibrium is defined by the equilibrium constant (Keq), representing the ratio of products to reactants.
• In cells, reactions are in a steady state, meaning intermediates are formed and consumed at the same rate.
• Homeostasis refers to the stable conditions maintained within cells due to this steady state.

pH and Buffers

• pH is the negative logarithm of hydrogen ion concentration, with cytoplasmic pH around 7.2-7.4.
• Acids release hydrogen ions, while bases bind them.
• Buffers are mixtures of weak acids and their conjugate bases that minimize pH changes.

Biochemical Energetics

• Energy types include kinetic energy (energy of motion), potential energy (stored energy), concentration gradients, and electric potential.
• Free energy (ΔG) measures the energy available for a reaction.
• Spontaneous reactions have negative ΔG (exergonic).
• Catalysts, like enzymes, lower activation energy to speed reactions without being consumed.
• ATP is the universal energy currency in cells, storing energy in phosphoanhydride bonds that release energy upon hydrolysis.
• Redox reactions involve electron transfer: oxidation (loss of electrons) and reduction (gain of electrons); electron carriers shuttle electrons.

Additional Notes

• The flow of genetic information is unidirectional (DNA → RNA → Protein).
• This flow controls cellular function and gene expression.