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Chemical Level of Organization and Biomolecules

Atoms and Molecules

  • Levels of structural organization: six levels (context for body organization).
  • Matter: occupies space and has mass; Mass = amount of material; On Earth, mass equals weight.
  • Atoms: smallest stable units of matter; contain protons (p+), neutrons (n0), and electrons (e-).
  • Subatomic particles:
    • Protons: positive charge
      n- Neutrons: neutral
    • Electrons: negative charge; much smaller and ~1/1800 the mass of protons/neutrons.
  • Atomic number (Z): number of protons; Mass number (A): protons + neutrons.
  • Isotopes: same element, different number of neutrons; differ in mass.
  • Atomic weight (amu): actual mass of an atom; expressed in daltons (amu).
  • Example: Hydrogen
    • Z = 1, A = 1, atomic weight ≈ 1.0079
  • Mole concept:
    • A mole (mol) is a quantity of matter; mass in grams equals an element’s atomic weight.
    • 1 mole of O = 16 ext{ g}; 1 mole of H = 1 ext{ g}.
  • Molecular weight: sum of atomic weights in a molecule; used to calculate reactant/product quantities.

Elements, Ions, and Bonds

  • Principal elements: thirteen most abundant elements by body weight; symbols include O, C, H, N, Ca, P, K, Na, Cl, Mg, S, I, etc.
  • Element symbols: most derived from English names; some from other languages (e.g., Na for sodium from natrium).
  • Major concept: electrons occupy energy levels (electron shells) around the nucleus.
  • Reactive vs inert elements:
    • Reactive: have unfilled outer energy levels (e.g., H); tend to gain/lose/share electrons.
    • Inert: outer energy levels filled (e.g., He, Ne); do not react easily.
  • Ionic bonds: formed by transfer of electrons (cations and anions attract).
    • Example formation: Na → Na⁺, Cl → Cl⁻; Na⁺ and Cl⁻ attract to form NaCl.
  • Covalent bonds: sharing of electrons between atoms.
    • Nonpolar covalent: electrons shared equally (no partial charges).
    • Polar covalent: unequal sharing; partial charges (e.g., H₂O: H slightly positive, O slightly negative).
  • Key examples:
    • H₂, O₂, CO₂ (covalent bonds; often polar in H₂O context)
    • H₂O (polar covalent bonds)
  • Electron transfer and bond stability lead to common bond types:
    • Ionic bonds: electrical attraction between ions
    • Covalent bonds: sharing electrons
  • Hydrogen bonds: weak attractions between polar molecules (e.g., between water molecules).

Chemical Reactions and Energetics

  • Cells stay alive by regulating chemical reactions: bonds form/break to form products.
  • Reactants → products; metabolism encompasses all cellular reactions.
  • Work and energy concepts:
    • Kinetic energy: energy of motion
    • Potential energy: stored energy
  • Chemical reactions notation:
    • Reactants on the left; products on the right. Example: 2H + O
      ightarrow H_2O
  • Types of chemical reactions (overview):
    • Decomposition: AB → A + B
    • Hydrolysis: AB + H₂O → AH + BH (insertion of water)
    • Synthesis: A + B → AB
    • Dehydration synthesis (condensation): A–H + OH–B → AB + H₂O
    • Exchange: AB + CD → AD + CB
  • Reversibility and equilibrium: many reactions are reversible; at equilibrium, forward and reverse rates are balanced.
  • Activation energy: energy required to start a reaction; enzymes lower activation energy to accelerate reactions.

The Importance of Water in the Body

  • Water is the most important body constituent; ~2/3 of body weight.
  • Properties:
    • Lubrication: reduces friction between surfaces (joints, cavities)
    • Reactivity: participates in hydrolysis and dehydration synthesis
    • High heat capacity: absorbs/retains heat; slows temperature changes (thermal inertia)
    • Solubility: many inorganic and organic molecules dissolve in water; water = solvent; solutes dispersed in solution
  • Hydrophilic vs hydrophobic: polar covalent (water-loving) vs nonpolar (water-fearing)
  • Ions and electrolytes: ions in solution conduct electricity (e.g., Na⁺, Cl⁻); ion concentrations are tightly regulated
  • Hydration, colloids, and suspensions explained: large organic molecules in solution; colloids stay in solution; suspensions may settle out
  • Water in chemical reactions: essential for hydrolysis and dehydration synthesis

pH and Buffers

  • Hydrogen ion (H⁺) and hydroxide ion (OH⁻) definitions; water dissociation yields both.
  • pH: negative logarithm of hydrogen ion concentration; ext{pH} = -\, ext{log}_{10}([ ext{H}^+]); ranges 0–14.
  • Blood pH: normally 7.35 ext{ to } 7.45.
    • Acidosis: pH < 7.35
    • Alkalosis: pH > 7.45
  • Acids and bases:
    • Acid: dissociates to release H⁺ (e.g., HCl → H⁺ + Cl⁻)
    • Base: removes H⁺ from solution (e.g., NaOH → Na⁺ + OH⁻)
  • Salts: ionic compounds made of cations other than H⁺ and anions other than OH⁻ (e.g., NaCl → Na⁺ + Cl⁻).
  • Buffers: stabilize pH by removing or releasing H⁺; usually involve a weak acid and its salt (e.g., H₂CO₃ / NaHCO₃).

Organic Compounds

  • Organic compounds contain carbon and hydrogen (often oxygen) and form long covalently bonded chains.
  • Generally soluble in water; functional groups influence properties and reactivity.

Carbohydrates

  • Formula ratio: $$C:H:O \