BIO 111 • Chapter 2A Study Notes – Atoms, Small Molecules, and Water

Atoms and Elements

• Atom vs. Element- Atom: Smallest unit of matter retaining element properties; made of protons, neutrons, electrons.

  • Element: Pure substance of one atom type; cannot be chemically simplified.

• Essential Elements- Definition: Elements needed in large amounts for survival, growth, reproduction (e.g., C, H, O, N, P, S, Ca, K, Na, Cl, Mg, Fe).

  • Periodic‐Table Location: First four periods; Groups 1–3 (metals: Na, K, Ca, Mg) and Groups 14–17 (non-metals: C, N, O, P, S, Cl).

• Elements Making Up Most Living Matter- Four primary atoms contribute
  \approx96 \%
  of mass in most organisms: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N).

Subatomic Particles

• Proton (p⁺)
  

  • Charge: +1
  

  • Mass: \approx1\,\text{amu} (atomic mass unit).
  

  • Location: Atomic nucleus.

• Neutron (n⁰)
  

  • Charge: 0
  

  • Mass: \approx1\,\text{amu} (slightly heavier than a proton).
  

  • Location: Nucleus.

• Electron (e⁻)
  

  • Charge: -1
  

  • Mass: \approx\tfrac{1}{1840}\,\text{amu} (negligible).
  

  • Location: Orbitals/electron cloud surrounding nucleus.
  

  • Chemical-reaction role: Electrons—especially those in the valence shell—are directly involved in bond formation and breaking.

Atomic Number, Mass, & Particle Count

• Atomic Number (Z): Number of protons. Defines the element.

• Mass Number (A): Protons + neutrons. A = p^+ + n^0

• Determining Particles (neutral atom):
  

  • #\,e^{-}=#\,p^{+}=Z
  

  • #\,n^{0}=A-Z

• Why Table Lists Whole Z but Decimal Mass: Z is integer; atomic mass is weighted average of isotopes.

Isotopes vs. Ions

• Isotopes: Same Z (element), different A (neutrons); neutral charge; altered mass; identical chemical behavior; some radioactive.

• Ions: Atom/molecule with net charge (lost/gained e^-). Cation: positive (lost e^-). Anion: negative (gained e^-). Altered charge affects electrical, osmotic, chemical properties.

Electron Shells & Reactivity

• Electron Shell (Energy Level): Region where electrons with similar energy are found.

• Filling Order (first 3 shells) 1. 1st shell: Max 2 e^-.

  2. 2nd shell: Max 8 e^-.

  3. 3rd shell: Max 8 e^- (for main-group elements). Lower energy shells fill first.

• Valence Shell: Outermost occupied shell; valence electrons determine reactivity.

• Inert vs. Reactive
  

  • Inert: Full valence shell (e.g., noble gases); no tendency to gain/lose/share e^-.
  • Reactive: Incomplete valence shell; seeks stability via bond formation (Octet Rule: 8 valence e^- optimal, 2 for H & He).

• Atoms ‘Seek’: Minimization of potential energy by achieving full/empty valence shells through bonding.

Electronegativity & Polarity

• Electronegativity (EN): Atom’s tendency to attract shared electrons in a covalent bond.

• Approximate Trend (Pauling scale): O (3.5) > N (3.0) > C (2.5)
  \approx
  H (2.1).

• Polarity: Unequal sharing of electrons (differing EN) → partial charges (δ⁺, δ⁻).

Types of Chemical Bonds

• Ionic Bonds
  

  • Complete e^- transfer (metal to non-metal); electrostatic attraction; medium strength, weaker in water.

• Covalent Bonds
  

  • Shared e^- pair(s). Non-polar: Equal sharing (ΔEN < 0.4) e.g., \text{C–H}, \text{O}2. Polar: Unequal sharing (ΔEN 0.5–1.9) → partial charges e.g., \text{H–O} in H2O. Strongest biological bonds.

• Hydrogen Bonds (H-bonds)
  

  • Electrostatic attraction between δ⁺ H (bonded to O, N, F) and δ⁻ EN atom of another molecule. Individually weak but numerous → major stabilizing force (DNA, proteins, water).

Valence Electrons, Vacancies, & Bonding Capacity

Formulas vs. Equations

• Molecular/Chemical Formula: Symbolic representation of composition (e.g., C6H{12}O_6).

• Chemical Equation: Depicts reaction (reactants → products), often balanced (e.g., 2H2 + O2 \rightarrow 2H_2O).

• Reactants: Consumed substances. Products: Formed substances.

Solutions & Solubility

• Solution: Homogeneous mixture (Solvent: greatest amount; Solute: dissolved components).

• Hydrophilic vs. Hydrophobic
  

  • Hydrophilic: Polar/charged; form H-bonds/ionic interactions with H_2O → soluble.
  • Hydrophobic: Non-polar; cannot interact with water → aggregate, insoluble, drive membrane formation.
  • Polarity dictates solubility/reactivity.

Water: Molecular & Intermolecular Bonds

• Intramolecular: Two polar covalent \text{H–O} bonds within each H_2O molecule.

• Intermolecular: Hydrogen bonds between δ⁺ H of one molecule and lone-pair O of another.

Cohesion vs. Adhesion

• Cohesion: Attraction between water molecules via H-bonds (surface tension, water columns).

• Adhesion: Attraction between water and other polar surfaces (capillary rise).

Water as ‘Universal Solvent’

• Polarity + up to 4 H-bonds → surrounds ions/polar molecules (hydration shells) and dissociates ionic lattices.

Thermal Concepts

• Kinetic Energy (KE): Energy of motion; for molecules, proportional to \frac{1}{2}mv^2.

• Temperature: Average KE of particles. Higher temp = faster motion.

• Specific Heat of Water: High because many H-bonds must break. c_{\text{water}} \approx 1\,\text{cal·g}^{-1}·°C^{-1} (about $2\text{x}$ typical liquids); moderates climate/organismal temp.

Density Anomaly of Ice

• Rigid lattice maximizes H-bond spacing → molecules fixed farther apart → lower density.

• Biological Significance
  

  • Ice floats, insulating water bodies → aquatic life survives; seasonal turnover oxygenates lakes.