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• Biochemistry evolved from organic chemistry and physiology to explain life processes at the molecular level.
• Ancient Indian medicine recognized metabolic basis of disease — Charaka described diabetes (madhumeha) as a disorder of carbohydrate and fat metabolism.
• Term “Biochemistry” coined by Neuberg in 1903 (bios = life, chymos = juice).
• Early foundational books:
– Liebig (1842): introduced concept of metabolism.
– Hoppe-Seyler (1877): published Textbook of Physiological Chemistry.
• Key early discoveries:
– Rouelle (1773): isolated urea.
– Wöhler (1828): synthesized urea → disproved “vital force.”
– Pasteur (1860): fermentation as biological process.
– Buchner (1897): extracted enzymes → cell-free fermentation.
• Major biochemical milestones:
– ATP isolated (Fiske & Subbarow, 1926).
– Creatine phosphate discovered (Lohmann, 1932).
– TCA cycle described (Krebs, 1937).
– DNA proved genetic material (Avery & MacLeod, 1944).
– DNA structure elucidated (Watson & Crick, 1953).
– Genetic code cracked (Nirenberg, 1961).
– Gene synthesis (Khorana, 1965).
• Modern breakthroughs:
– Recombinant DNA (Paul Berg, 1972).
– PCR (Kary Mullis, 1985).
– Human Genome Project (1990–2003).
• Present-day importance:
– Central to understanding metabolism, genetics, enzymology, disease mechanisms.
– Basis of modern diagnostics, biomarkers, molecular medicine, and therapeutics.
• Six elements form >99% of body composition: O, C, H, N, Ca, P.
• Body composition: 60% water, 15% proteins, 15% lipids, 2% carbohydrates, 8% minerals.
• All major biomolecules arise from ~30 basic precursors (“alphabet of biochemistry”).
• Includes 20 amino acids, 2 purines, 3 pyrimidines, glucose, ribose, palmitate, glycerol, choline.
• Monomers → Macromolecules → Supramolecular assemblies → Organelles → Cell → Tissue.
• Examples:
– Glucose → glycogen
– Amino acids → proteins
– Nucleotides → nucleic acids
• Biomolecules join by covalent bonds (strong, stable).
• Examples:
– Glycosidic bonds (carbohydrates)
– Peptide bonds (proteins)
– Phosphodiester bonds (nucleic acids)
• Covalent assembly forms large macromolecules essential for structure and function.
• Macromolecules interact via noncovalent forces (hydrogen bonds, ionic bonds, van der Waals, hydrophobic interactions).
• Examples:
– Ribosomes (RNA + proteins)
– Lipoproteins (lipids + proteins)
– Chromatin (DNA + protein)
• Prokaryotes:
– Small size
– No true nucleus
– No membrane-bound organelles
– Rigid cell wall
• Eukaryotes:
– Large size (1000–10,000× larger)
– True nucleus with nuclear membrane
– Organelles present (mitochondria, ER, Golgi, lysosomes)
– Plasma membrane with lipid bilayer
• Dietary carbohydrates, fats, and proteins are broken down into monosaccharides, fatty acids/glycerol, and amino acids.
• Takes place in the gastrointestinal tract by enzymatic hydrolysis.
• Produces small absorbable molecules that enter circulation.
• Absorbed nutrients undergo further oxidation inside cells.
• Generates NADH and FADH₂.
• Includes pathways like glycolysis, β-oxidation, deamination, and TCA cycle.
• Main purpose: extraction of reducing equivalents and metabolic intermediates.
• NADH and FADH₂ feed electrons to the electron transport chain in mitochondria.
• Electron flow creates a proton gradient used by ATP synthase to form ATP.
• Oxygen acts as the final electron acceptor, producing water.
• Catabolism: Breakdown of complex molecules, energy-releasing, exergonic (e.g., glycolysis, β-oxidation).
• Anabolism: Synthesis of biomolecules, energy-requiring, endergonic (e.g., glycogen, protein, and fatty acid synthesis).
• Together they maintain overall metabolic balance in the body.
• Strong bonds formed by sharing of electron pairs between atoms.
• Responsible for basic structure of biomolecules (peptide, glycosidic, phosphodiester bonds).
• High bond strength → stable macromolecular backbones.
• Formed by attraction between oppositely charged ions.
• Occur when an electron is transferred from an electropositive atom to an electronegative atom.
• Strength decreases in aqueous environments but is important for protein folding and interactions.
• Na⁺–Cl⁻, K⁺–Br⁻, Na⁺–F⁻ represent classical ionic compounds.
• Demonstrate electron transfer and electrostatic attraction between the resulting ions.
• Positive (basic) groups:
– Lysine: ε-amino group
– Arginine: guanidinium group
– Histidine: imidazolium group
• Negative (acidic) groups:
– Aspartate: β-carboxyl group
– Glutamate: γ-carboxyl group
• These charged groups form ionic interactions that stabilize protein tertiary and quaternary structure.
• Atoms or groups that supply a hydrogen already covalently attached to an electronegative atom.
• Common donors in biological molecules include:
– –NH groups (as in peptide bonds, indole ring of tryptophan, imidazole ring of histidine)
– –OH groups (serine, threonine, tyrosine)
– –NH₂ groups (lysine, arginine side chains)
• Electronegative atoms with lone pairs of electrons capable of accepting a hydrogen.
• Major acceptors in biomolecules:
– Carbonyl oxygen (C=O) of peptide bonds
– Carboxylate groups (COO⁻) of aspartate and glutamate
– Sulfur atoms in disulfide linkages may participate weakly
– Nitrogen or oxygen atoms in nucleic acid bases
• In proteins:
– Hydrogen bonds stabilize α-helices, β-sheets, and turns of secondary structure.
– Occur between the carbonyl oxygen of one amino acid and the amide hydrogen of another along the peptide backbone.
– Side-chain hydrogen bonds help fine-tune folding and maintain the tertiary structure.
• In DNA:
– Hydrogen bonds hold complementary bases together:
– A–T pairs form two hydrogen bonds.
– G–C pairs form three hydrogen bonds.
– These bonds allow the double helix to be stable yet separable, enabling replication and transcription.
– Provide specificity in base pairing and maintain the geometry of the helix.
• Nonpolar amino-acid side chains (like leucine, isoleucine, valine, phenylalanine) tend to avoid water.
• In an aqueous environment, these nonpolar groups move closer together, reducing their contact with water.
• This clustering forms the hydrophobic core of proteins and is a major driving force behind protein folding.
Nonpolar side chains come together and pack tightly in the interior of a protein while water molecules remain on the outside. This minimizes the exposure of hydrophobic residues to the surrounding water, helping the protein fold into a stable three-dimensional structure.
• Very weak, short-range attractive forces arising from temporary dipoles formed when electrons fluctuate around atoms.
• Occur between all atoms, whether polar or nonpolar.
• Individually weak (≈1 kcal/mol), but collectively significant because they act throughout the protein interior.
• Essential for tight packing of amino acids within the folded structure and stabilizing the nonpolar core.
• Each water molecule can form up to four hydrogen bonds due to its bent structure and polarity.
• These dynamic bonds constantly form and break, creating a flexible network in liquid water.
• This network gives water its unique physical and chemical properties essential for life.
• In ice, water molecules arrange in a fixed tetrahedral structure, creating more open space.
• This makes ice less dense than liquid water, allowing it to float.
• When ice melts, hydrogen bonds collapse partially, and molecules pack more tightly → higher density.
• As temperature increases, water molecules move faster, breaking hydrogen bonds more frequently.
• At 0–4°C, collapsing hydrogen bonds increase density.
• Above 4°C, thermal expansion dominates → density decreases.
• At 100°C, kinetic energy exceeds hydrogen bonding → water becomes vapor.
• Hydrophilic molecules (polar or charged) dissolve easily because they can form hydrogen bonds or ionic interactions with water.
• Hydrophobic molecules (nonpolar) cannot form bonds with water, causing them to cluster together, a key principle in protein folding and membrane formation.
Water molecules are shown connected through multiple hydrogen bonds: the oxygen atom of one molecule attracts the hydrogen atom of a neighboring molecule. These repeated interactions create a loosely organized yet constantly shifting network that explains water’s high cohesion, surface tension, and solvent capacity.
• Total energy of a system and its surroundings remains constant.
• Energy can be transformed but not created or destroyed.
• ΔE: Change in internal energy
• Q: Heat absorbed by the system
• W: Work done by the system
• If heat is absorbed and no work is done, internal energy increases.
• Spontaneous processes increase the entropy (randomness/disorder) of a system + surroundings.
• Systems naturally move from ordered → disordered states unless energy is added.
• Entropy (S) reflects randomness; higher entropy means greater disorder.
• At equilibrium, entropy reaches its maximum.
• Biological systems maintain order by increasing entropy in surroundings.
• Q: Heat absorbed
• T: Absolute temperature
• ΔS: Change in entropy
• Shows that entropy changes depend on heat flow and temperature.
• Determines whether a reaction can occur spontaneously.
• Combines enthalpy (heat) and entropy (disorder) effects.
• ΔG: Free energy change
• ΔH: Change in enthalpy
• TΔS: Temperature × entropy change
• Negative ΔG → reaction proceeds spontaneously.
• Positive ΔG → requires energy input.
• Exergonic: ΔG negative → energy-releasing → spontaneous.
• Endergonic: ΔG positive → energy-requiring → non-spontaneous unless coupled.
• Unfavorable reactions (positive ΔG) can proceed when paired with favorable reactions (negative ΔG).
• The combined ΔG becomes negative, allowing the pathway to move forward.
• Glucose + Pi → Glucose-6-phosphate is endergonic (positive ΔG).
• Coupled with ATP hydrolysis (strongly exergonic).
• Net ΔG becomes negative, enabling phosphorylation of glucose.
• Free energy change under standard conditions (1 M concentration, pH 7).
• Indicates reaction tendency but actual ΔG depends on concentrations inside cells.
• Near-equilibrium reactions:
– ΔG ≈ 0
– Easily reversible
– Direction depends on substrate/product concentration.
• Far-from-equilibrium reactions:
– Large negative ΔG
– Irreversible
– Usually catalyzed by key regulatory enzymes.
• Pathways contain a few strongly exergonic steps that drive the entire sequence forward.
• These irreversible steps determine pathway direction, regulate metabolic flow, and prevent backward cycling.
• When two solutions are separated by a membrane permeable to water and small ions but not to large charged molecules (e.g., proteins), ions redistribute unevenly.
• The side containing non-diffusible ions develops a predictable imbalance of diffusible ions.
• At equilibrium, the product of concentrations of diffusible cations and anions on one side equals the product on the other side.
• This rule explains why ion distribution becomes asymmetric when large fixed ions are present.
• Each compartment must maintain overall neutrality:
– Total cations = total anions.
• Even if ions shift across the membrane, each side still balances charge internally.
• The total amount of each diffusible ion (e.g., Na⁺ or Cl⁻) remains equal before and after equilibrium.
• Only the distribution between the two sides changes.
One side of the membrane contains a salt with a non-diffusible anion (such as a protein), while the other side contains only diffusible ions. As equilibrium is reached, more cations accumulate on the side with the non-diffusible anion, and fewer anions remain there. The other side shows the complementary pattern. Despite this imbalance, the equations for neutrality, product rule, and conservation all remain satisfied.
• Plasma proteins (non-diffusible anions) cause an uneven distribution of small ions.
• This contributes to the oncotic pressure that helps maintain plasma volume.
• Cells contain negatively charged proteins → attract more H⁺ ions.
• This leads to slightly lower pH inside tissues compared to extracellular fluid.
• Hemoglobin carries strong negative charges.
• This causes increased H⁺ concentration inside red cells → lower RBC pH than plasma.
• When CO₂ enters RBCs and is converted to HCO₃⁻, chloride ions move into RBCs to maintain electrical neutrality.
• This unequal movement of Cl⁻ is partly explained by the Donnan effect.
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