Principles of Life: Chemistry of Life

Matter and Elements

  • Matter = anything that has mass and occupies space.
  • Elements cannot be broken down into simpler substances.
  • There are 92 natural elements.
  • Living organisms are made primarily of six elements: C, H, O, N, P, S.

The Human Body: Major Elements and Abundance

  • Major elements by weight in the human body (approx.):
    • Oxygen: 65.0\%
    • Carbon: 18.5\%
    • Hydrogen: 9.6\%
    • Nitrogen: 3.3\%
    • Calcium: 1.5\%
    • Phosphorus: 1.0\%
    • Potassium: 0.4\%
    • Sulfur: 0.3\%
    • Sodium: 0.2\%
    • Chlorine: 0.2\%
    • Magnesium: 0.1\%
  • Trace elements (each < 0.01%): B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn.

Atomic Structure and the Periodic Table

  • An atom consists of:
    • Nucleus: protons (positive) and neutrons (neutral).
    • Electrons: move around the nucleus.
  • Atomic number = number of protons; defines element.
  • Atomic weight = weighted average of isotopes' masses.
  • Periodic table: elements in the same vertical group have the same number of electrons in their outermost shell (valence electrons).

Atoms, Molecules, and Bonds

  • A molecule is two or more atoms bonded together; a compound contains two or more different elements.
  • Bond types hold molecules together; bonds influence stability and reactivity.
  • Valence electrons determine bonding behavior (outer-shell electrons).

Chemical Bonds: Why Bonds Form

  • Atoms seek a stable (low-energy) state by filling their valence shell (octet rule: eight valence electrons).
  • Bonds form via gaining, losing, or sharing electrons.

Covalent Bonds

  • Covalent bonds: atoms share electrons to complete valence shells.
  • Covalent bonds are very strong.
  • Example: two hydrogen atoms form H₂ via a covalent bond.
  • Carbon forms covalent bonds and can create long chains and rings (tetravalent): single, double, or triple bonds.

Representations of Molecules

  • Common representations include:
    • Electron distribution diagrams
    • Structural formulas
    • Molecular models
  • Examples of bond types:
    • H₂: single bond
    • O₂: double bond
    • CH₄: methane (single bonds)
    • C₂H₄: ethene (double bonds)
    • C₂H₂: acetylene (triple bonds)
    • Aromatic rings like benzene (C‑H connections in rings)

Properties of Covalent Bonds

  • Orientation is fixed: bond length, angle, and direction are consistent.
  • Strength and stability: covalent bonds are strong; bond strengths affect reaction likelihood.

Polar Covalent Bonds and Electronegativity

  • When atoms have different electronegativities, sharing is unequal.
  • More electronegative atom becomes δ−; the other becomes δ+.
  • Example: water (O–H) exhibits polarity.

Ionic Bonds

  • Ionic bonds form when electrons are transferred, creating oppositely charged ions.
  • Salts (e.g., NaCl) consist of ions held by ionic bonds; they often dissociate in water.
  • Ions (electrolytes) are important for nerve impulses, muscle contractions, and water balance.
  • In NaCl, Na tends to lose an electron (Na⁺) and Cl tends to gain an electron (Cl⁻).

Lattice Structure of Ionic Compounds

  • Ionic compounds form lattice structures (e.g., NaCl) with alternating Na⁺ and Cl⁻ ions.

Hydrogen Bonds and van der Waals Forces

  • Hydrogen bonds: attractions between oppositely charged parts of molecules (weakest of the main bonds).
  • Important in DNA structure and water properties (surface tension).
  • van der Waals attractions: transient, shape-dependent interactions between molecules.

Bonds in Cells

  • Strong covalent bonds build DNA backbones; hydrogen bonds stabilize DNA double helix.
  • Weak interactions allow different molecules and ions to interact dynamically (e.g., Na⁺ with water, then with a protein).

Water and Life

  • Water is the most important molecule for life; organisms are ~70–90% water.
  • Water properties driven by hydrogen bonding:
    • Adhesion and cohesion
    • Surface tension
  • Water as solvent: hydrophilic (water-loving) vs hydrophobic (water-fearing) substances.
  • Water dissolves polar, ionic, and some nonpolar gases.

Water in Action: Heat, Cooling, and Ice

  • Water has high heat capacity and high heat of vaporization, stabilizing temperatures in organisms.
  • Evaporative cooling: sweating uses water's high heat of vaporization to remove heat.
  • Frozen water is less dense than liquid water; ice expands when it freezes and floats.

Water and the Search for Life

  • Liquid water is a key criterion in the search for extraterrestrial life; NASA recognizes its importance for life as we know it.

Acids, Bases, and pH

  • In water, dissociation yields H⁺ and OH⁻ ions:
    \mathrm{H_2O \rightleftharpoons H^+ + OH^-}
  • Acids release H⁺ in solution (high H⁺ concentration).
  • Bases release OH⁻ in solution (low H⁺ concentration).
  • Example acids/bases:
    • Hydrochloric acid: \mathrm{HCl \rightarrow H^+ + Cl^-}
    • Sodium hydroxide: \mathrm{NaOH \rightarrow Na^+ + OH^-}

pH and Buffers

  • pH measures the concentration of H⁺ in solution; lower pH = more acidic, higher pH = more basic.
  • Buffers resist changes in pH by absorbing or releasing H⁺/OH⁻ as needed.

Organic vs Inorganic Chemistry

  • Organic chemistry = chemistry of living world; inorganic chemistry = chemistry of nonliving world.
  • An organic molecule must contain carbon and hydrogen.

Chemical Reactions: Reactants and Products

  • A chemical reaction converts reactants to products.
  • Photosynthesis example: 6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2