Chemistry Essentials for Biology (Transcript Notes)

Chemistry foundations for biology, atomic structure, essential and trace elements, isotopes and radioactivity, chemical bonding, water’s properties, electrolytes, and quick review questions. Includes key formulas in LaTeX and example scenarios mentioned in the lecture.

1. Chemistry in Biology

  • Macromolecules: Understanding the chemical composition and structure of carbohydrates, lipids, proteins, and nucleic acids is fundamental to biology.
  • Metabolism: Chemical reactions drive all biological processes, including energy production (e.g., cellular respiration) and synthesis of biomolecules.
  • Interactions: The specificity of interactions between molecules (e.g., enzyme-substrate binding, hormone-receptor interaction) is governed by chemical principles.

2. Atomic Structure

  • Subatomic Particles:
    • Protons (p^+$): Positively charged, located in the nucleus. Determine the atomic number (Z) and identity of an element.
    • Neutrons (n^0): No charge, located in the nucleus. Contribute to atomic mass, along with protons.
    • Electrons (e^ -): Negatively charged, orbit the nucleus in electron shells or orbitals. Determine the chemical properties and reactivity of an atom.
  • Atomic Number (Z): Number of protons in an atom's nucleus. Unique to each element.
  • Mass Number (A): Total number of protons and neutrons in an atom's nucleus. Expressed as Mass ext{ }Number = Protons + Neutrons.
  • Atomic Weight: The weighted average of the mass numbers of all naturally occurring isotopes of an element, measured in atomic mass units (amu).

3. Essential and Trace Elements

  • Major Essential Elements: Elements that make up the bulk of living matter.
    • CHNOPS: Carbon (C), Hydrogen (H), Nitrogen (N), Oxygen (O), Phosphorus (P), and Sulfur (S). These six elements comprise approximately 96% of living organisms.
    • Importance: They form the backbone of organic molecules and are crucial for all biological structures and functions.
  • Trace Elements: Elements required in very small quantities for normal growth and development.
    • Examples: Iron (Fe), Copper (Cu), Zinc (Zn), Iodine (I).
    • Role: Often serve as cofactors for enzymes or are components of essential biomolecules (e.g., Iron in hemoglobin).

4. Isotopes and Radioactivity

  • Isotopes: Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers.
    • Example: Carbon-12 ($^{12}C), Carbon-13 ($^{13}C), and Carbon-14 ($^{14}C).
  • Radioactive Isotopes (Radioisotopes): Unstable isotopes whose nucleus decays spontaneously, releasing energy and particles (radiation).
    • Half-life: The time it takes for half of the radioactive atoms in a sample to decay. Each radioisotope has a characteristic half-life (T_{1/2}).
    • Applications:
      • Medical Imaging: PET scans use radioisotopes diagnostically.
      • Carbon Dating: {}^{14}C is used to date organic materials up to tens of thousands of years old.
      • Tracers: Used to follow metabolic pathways in research.
    • Risks: Radiation can damage DNA, leading to mutations or cancer.

5. Chemical Bonding

  • Valence Electrons: Electrons in the outermost shell, determining an atom's reactivity and ability to form bonds.
  • The Octet Rule: Most atoms strive to achieve 8 valence electrons (or 2 for small atoms like H and He) to attain stability.
  • Types of Bonds:
    • Ionic Bonds: Formed by the complete transfer of electrons between atoms, resulting in the formation of ions (cations and anions) that are attracted by electrostatic forces.
      • Example: NaCl (Sodium chloride) -> Na^+ and Cl^-.
    • Covalent Bonds: Formed by the sharing of electrons between atoms.
      • Nonpolar Covalent Bond: Equal sharing of electrons between atoms with similar electronegativity (e.g., $O2$, $CH4).
      • Polar Covalent Bond: Unequal sharing of electrons between atoms with different electronegativity, creating partial positive ( ext{δ}+) and partial negative ( ext{δ}-) charges (e.g., H_2O).
    • Hydrogen Bonds: Weak attractions that occur when a hydrogen atom covalently bonded to a highly electronegative atom (like O, N, or F) is attracted to another highly electronegative atom in a different molecule or part of the same molecule.
      • Crucial for water's properties and protein/DNA structure.
  • Electronegativity: An atom’s attraction for shared electrons in a covalent bond.

6. Water’s Properties

  • Polarity: Due to its bent shape and the high electronegativity of oxygen, water (H_2O) is a highly polar molecule, allowing it to form extensive hydrogen bonds.
  • Cohesion: Water molecules stick to each other via hydrogen bonds, leading to high surface tension.
  • Adhesion: Water molecules stick to other polar surfaces.
  • High Specific Heat: Water can absorb or release a large amount of heat with only a slight change in its own temperature, moderating climate and regulating organismal temperature.
    • Specific Heat of Water: 4.184 ext{ }J/(g ext{ }°C).
  • High Heat of Vaporization: A large amount of energy is required to convert liquid water into water vapor, enabling evaporative cooling.
  • Versatile Solvent: Water's polarity allows it to dissolve many ionic and polar substances, making it the “universal solvent” for biological processes. Hydrophilic substances dissolve in water; hydrophobic substances do not.

7. Electrolytes

  • Definition: Substances that dissociate into ions when dissolved in water, making the solution capable of conducting electricity.
  • Biological Importance: Essential for nerve impulse transmission, muscle contraction, fluid balance (osmosis), and maintaining proper pH levels.
  • Examples: Ions like sodium (Na^+), potassium (K^+), chloride (Cl^-$), calcium