In-depth Chemistry for Cell Biology Notes

Introduction to Cell Biology Chemistry

  • Importance of fundamental chemistry knowledge for biology
    • Key concepts are necessary for all students, regardless of prior experience.
    • Related material available in Karp textbook (Chapter 2, sections 2.1-2.4).

Composition of Typical Eukaryotic Cells

  • Major classes of molecules (approximate amounts):
    • Protein: 10%
    • Nucleic Acid: 1.1%
    • DNA: 0.4%, RNA: 0.7%
    • Lipid: 0.2%
    • Fats, oils, steroids
    • Carbohydrate: 0.4%
    • Sugars, starches
    • Inorganic matter: 1.5%
    • Calcium (Ca), Magnesium (Mg)
    • Water: 85%

Key Atoms in Cell Biology

  • Six essential atoms: CHNOPS

    • C: Carbon
    • H: Hydrogen
    • N: Nitrogen
    • O: Oxygen
    • P: Phosphorus
    • S: Sulfur

  • Arrangement of these atoms in biomolecules influences their formation:

    • Proteins, nucleic acids, lipids, carbohydrates structured from these atoms.
    • Important bonding rules:
    • Carbon: 4 bonds
    • Hydrogen: 1 bond
    • Nitrogen: 3 bonds
    • Oxygen: 2 bonds
    • Phosphorus: 5 bonds
    • Sulfur: 2 bonds

Atomic Structure

  • Composition of atoms: Nucleus + Electrons
    • Nucleus contains protons (positive charge) and neutrons (no charge).
    • Electrons (negative charge) exist in shells around nucleus.
    • Protons and electrons usually equalize for a net charge of 0.
Electron Shells and Bonding
  • Importance of outermost electrons for chemical bonding:
    • Example configurations:
    • Carbon (C): 2 in first shell, 4 in second
    • Oxygen (O): 2 in first, 6 in second
    • Sodium (Na): 2 in first, 8 in second, 1 in third
    • Chlorine (Cl): 2 in first, 8 in second, 7 in third

Chemical Bonds

  1. Covalent Bonds

    • Strongest type of bond, through sharing of electron pairs.
    • Single bond: one pair shared (example: H-H).
    • Double bond: two pairs shared (example: O=C=O).
    • Triple bonds: less common in biology.
    • Importance: affects the 3D structure, does not allow rotation around a double bond.

  2. Ionic Bonds

    • Formed by transfer of electrons.
    • Example: Sodium chloride (NaCl) - Na donates an electron to Cl.
    • Ions: Atoms with unequal proton/electron balance.
    • Ionic bonds dissociate in solution (e.g., Na+ and Cl- ions in water).

  3. Hydrogen Bonds

    • Weak bonds formed by attraction between polar molecules (e.g., water).
    • Essential for water's properties and biological interactions.
    • Polar vs non-polar covalent bonds:
      • Polar: electrons shared unequally, resulting in partial charges.
      • Non-polar: electrons shared equally.
    • Hydrophilicity (water-loving) vs hydrophobicity (water-fearing).

  4. Van der Waals Forces

    • Occur due to transient dipoles in nonpolar molecules.
    • Weak attractive forces contributing to molecular stability.

The Concept of pH

  • pH measures the concentration of hydrogen ions (H+).
  • Importance of pH in biological reactions:
    • Changes can significantly affect biological processes (e.g., acidosis).
  • Acid and base definitions:
    • Acids release H+ ions in solution (e.g., acetic acid: CH3COOH → CH3COO− + H+).
    • Bases reduce the concentration of H+ ions (e.g., ammonia: NH3 + H+ → NH4+).
pH Scale and Calculations
  • pH = -log10[H+]; pure water has a pH of 7.
  • Each unit change in pH corresponds to a tenfold change in H+ concentration (e.g., pH 5 is 100x more acidic than pH 7).
  • Understanding pH is crucial for maintaining homeostasis in biological systems.