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Section 2.1: Chemistry Foundations for Biology 3

Section 2.1: Chemistry Foundations for Biology

  • Purpose of this quick start:
    • Quick review of terminology to build a foundation for water, pH, inorganic molecules, and organic molecules to be covered later in the week.
    • Emphasis on model shifts in biology and how chemistry underpins cellular structures, digestion, and health.
    • Acknowledge that this subsection (2.1) will not be heavily tested, but terms will appear on tests later in the chapter.
  • Course logistics mentioned:
    • Tutor and past sessions available (e.g., 05:30–06:30). Check announcements.
    • If you’re new to chemistry, you must get up to date with terms to succeed across the course.
  • Core idea: Everything in biology rests on chemical elements, atoms, bonds, water properties, and pH; these concepts recur in carbohydrates, fats, proteins, nucleic acids, digestion (AMP1 and AMP2), and physiology.
  • Outcomes to know (highlights):
    • What is a chemical element? The simplest form of matter.
    • Elements are identified by their atomic number (the number of protons in the nucleus).
    • Elements are represented by one- or two-letter symbols on the periodic table; there are 91 elements.
    • Six elements are most abundant in humans: ext{O, C, H, N, Ca, P} (oxygen, carbon, hydrogen, nitrogen, calcium, phosphorus). Highlight these; they appear repeatedly in organic molecules and bones.
    • Nucleic acids contain nitrogen; bones involve calcium and phosphorus.
    • Trace elements exist in small amounts and can be important.
    • Some elements are minerals (inorganic) and lack carbon; others are organic (contain carbon).
    • Minerals of particular importance: ext{Ca}^{2+} ext{ and } ext{P} for bones, enzyme function, nerves, and muscles.
    • Atomic structure basics (not heavily tested here): protons = positive charge, neutrons = neutral, electrons = negative; nucleus contains protons and neutrons; electrons orbit the nucleus; atomic mass ≈ number of protons + neutrons; a neutral atom has equal numbers of protons and electrons; valence electrons are in the outer shell and drive bonding.
    • Isotopes: varieties of an element differing in neutron number; same number of electrons, but different atomic mass due to neutrons.
    • Radioisotopes are unstable isotopes that decay and emit radiation; potential hazards include ionizing radiation which can destroy molecules and generate free radicals.
    • Historical note: Marie (Madam) Curie discovered radioactivity and coined the term; polonium and radium were key elements; Curie died from radiation poisoning.
    • Ion vs ionization: an ion is a charged particle (atom or molecule) with unequal numbers of protons and electrons; ionization involves electron transfer.
    • Cation vs anion: cation has more protons (positive charge); anion has more electrons (negative charge).
    • Common example: sodium chloride (NaCl) forms via electron transfer from Na to Cl; Na becomes Na⁺, Cl becomes Cl⁻; salts are inorganic and typically dissociate in water; electrolytes are substances that ionize in water and conduct electricity; they contribute to osmotic balance and electrical excitability in nerves and muscles.
  • Basic chemistry foundations:
    • Atomic number Z = number of protons; determines identity of the element.
    • Atomic mass A ≈ Z + N (protons + neutrons), measured in atomic mass units (amu).
    • Isotopes differ in neutron number; radioisotopes are of particular interest due to radioactivity and ionizing radiation.
    • Ionization and electronic structure underpin chemical bonding and reactions.
  • Chemical bonds overview (to be applied to inorganic/organic molecules):
    • Ionic bond: transfer of electrons resulting in ongoing attraction between cation and anion (e.g., NaCl). Easily broken by water.
    • Covalent bond: atoms share electrons; can be:
    • Single covalent bond: one shared pair of electrons.
    • Double covalent bond: two shared pairs of electrons.
    • Bond polarity depends on how equally electrons are shared:
      • Nonpolar covalent bond: electrons shared equally.
      • Polar covalent bond: electrons shared unequally (slight partial charges).
    • Hydrogen bond: a weak attraction between a slightly positive hydrogen atom in one molecule and a slightly negative atom (usually O or N) in another molecule; crucial for water properties and biomolecules (DNA, proteins).
    • Van der Waals forces: brief attractions between neutral atoms caused by transient electron density fluctuations; weaker than covalent bonds (about ~1% strength of covalent bonds) but important in molecular interactions and protein folding.
  • Water and life: a critical focus area for the course
    • Water is polar and forms hydrogen bonds; this contributes to many key properties (solvent, cohesion, adhesion, chemical reactivity, thermal stability).
    • Water makes up ~60\% of the human body; its properties enable homeostasis and biochemical reactions.
    • Water as the universal solvent: dissolves more substances than any other solvent; essential for metabolic reactions.
    • Hydrophilic vs hydrophobic:
    • Hydrophilic substances dissolve in water (polar or charged).
    • Hydrophobic substances do not dissolve in water (nonpolar, e.g., fats and lipids).
    • Adhesion: water’s tendency to cling to other surfaces (membranes, serous membranes) and reduce friction around organs.
    • Cohesion: water molecules stick to each other, enabling surface tension and stable fluid movement.
    • Surface tension: cohesive forces at the air-water interface allow up to a spill-over beyond the rim in some cases.
    • Chemical reactivity: water participates in chemical reactions, can ionize into H⁺ and OH⁻; participates in hydrolysis and dehydration synthesis:
    • Hydrolysis: chemical reaction where water breaks a bond in a molecule (
      hydro = water; lysis = break) → break apart.
    • Dehydration synthesis: removal of water to form a bond and build molecules.
    • Thermal stability and high heat capacity: water requires a large amount of energy to change its temperature; this stabilizes body temperature and helps maintain homeostasis. Calorie as the base unit of heat for this context.
  • Mixtures and states of matter in biology:
    • Mixture: substances physically blended but not chemically bonded; each component retains properties.
    • Common body composition includes large amounts of water; mixtures in physiology include solutions, colloids, suspensions, and emulsions.
    • Key definitions (focus on solutions):
    • Solution: solute particles are small, dissolved in a solvent (usually water); transparent and pass through most membranes due to tiny particle size.
    • Colloid: larger particles than a solution; may scatter light and appear cloudy; typically do not pass through semipermeable membranes.
    • Suspension: very large particles (e.g., cells) that do not pass through membranes and appear cloudy.
    • Emulsion: suspension of one liquid in another (e.g., oil in water; can be emulsified with an emulsifier).
    • In biology, we mainly discuss solutions, but suspensions and emulsions are also relevant (e.g., blood plasma as a suspension of cells).
  • Practical health and clinical relevance:
    • Electrolytes (ions like Na⁺, K⁺, Ca²⁺) in water-based solutions are essential for osmotic balance and electrical excitability of nerves and muscles.
    • Age-related changes can impair the regulation of sodium and potassium, increasing risk for cardiac issues and overall health challenges.
    • Antioxidants (e.g., selenium, vitamins A and C, keratin) help neutralize free radicals; enzymes like superoxide dismutase (SOD) convert reactive species to less harmful products (e.g., superoxide to water and oxygen).
  • Organic vs inorganic molecules (context for later chapters):
    • Inorganic: typically lack carbon or do not contain carbon-hydrogen backbones in biomolecules.
    • Organic: contain carbon; form carbohydrates, fats, proteins, nucleic acids; building blocks and metabolic intermediates.
    • Emphasis on understanding simple definitions first, then applying them to complex biomolecules in digestion and metabolism.
  • Important numbers and quick references from the lecture:
    • Elements: 91 total; 6 most abundant in humans: ext{O}, ext{C}, ext{H}, ext{N}, ext{Ca}, ext{P}
    • Body composition: ~60\% water by mass.
    • pH scale basics (introduced later in detail): neutral pH is 7; acids have pH < 7; bases have pH > 7.
    • Blood pH target: approximately 7.4 (slightly basic) for homeostasis.
    • pH as a negative logarithm: \mathrm{pH} = -\log [\mathrm{H^+}]. A tenfold change in hydrogen ion concentration corresponds to a change of 1 pH unit on the scale:
    • If [\mathrm{H^+}] = 10^{-4}\,\text{M}, then \mathrm{pH} = 4; if [\mathrm{H^+}] = 10^{-5}\,\text{M}, then \mathrm{pH} = 5, etc.
  • Note on safety and ethics in radiation:
    • Ionizing radiation from radioisotopes can be hazardous (cancer, DNA damage) but is also used diagnostically and biologically in controlled ways.
    • Historical cautionary tale: exposure and radiation poisoning as highlighted by Madam Curie’s era; ongoing emphasis on safety and risk awareness in medical contexts.
  • Quick glossary (referenced terms):
    • Element, Atomic number, Atomic mass, Isotope, Radioisotope, Ion, Cation, Anion, Ionic bond, Covalent bond, Polar covalent, Nonpolar covalent, Hydrogen bond, Van der Waals forces, Ionization, Electrolyte, Salt, Mineral, Inorganic vs Organic, Solvent, Solute, Solution, Colloid, Suspension, Emulsion, Hydrophobic, Hydrophilic, Hydrolysis, Dehydration synthesis, pH, Acid, Base, Buffer.
  • How this connects to later topics:
    • The concepts of bonding, water properties, and pH underpin digestion (protein, carbohydrate, fat metabolism), cellular structure, and homeostasis.
    • Organic molecules (carbohydrates, lipids, proteins, nucleic acids) depend on covalent bonding and functional groups; water mediates reactions and structural stability.
    • Understanding minerals and electrolytes lays groundwork for physiology, nerve function, muscle contraction, and clinical care (e.g., electrolyte management in elderly patients).
  • Summary takeaway:
    • A solid grasp of elements, isotopes, ions, types of chemical bonds, water’s properties, and pH is essential for interpreting biology and medicine. These concepts recur across all subsequent modules (structure, digestion, physiology, health).