AP Biology – Chemistry of Life Review

Course & Assessment Logistics

• Instructor will skip introductory scientific-method material; assumes prior knowledge.

• No dissections in new AP Biology curriculum; focus is molecular → cell biology.

• Planned content sequence

  • Biomolecules (≈ 3 classes)

  • Cell structure & division (mitosis, meiosis)

  • Genetics

  • Metabolism (photosynthesis & cellular respiration)

  • Evolution

  • Ecology (self-study; slides provided, few formal lectures)

• Timeline

  • All instruction completed by December; mid-March unit testing begins.

  • AP exam now scheduled for first week of May; April reserved for school finals.

• Labs include DNA isolation; emphasis on safe, fun experiments.

Biology as a Multidisciplinary Science

• Sub-fields: molecular biology, genetics, cell biology, environmental science, ecology, paleontology, evolutionary biology, etc.

• Biologists study life under the constraints of physical laws: conservation of matter & energy, unidirectional energy flow, nutrient cycling.

Matter, Elements & Compounds

• Matter: has mass & occupies 3-D space (volume).

• Elements

  • Pure substances made of identical atoms; cannot be broken by chemical means.

  • Periodic table currently lists 118 known elements.

  • Essential elements for life (~25% of table)

  • 96 % of living mass: Carbon, Hydrogen, Oxygen, Nitrogen (CHON).

  • Remaining 4 %: Calcium, Phosphorus, Potassium, Sulfur.

  • Trace elements (needed in minute amounts)

  • Iron → hemoglobin; deficiency → anemia (common in females).

  • Iodine → thyroid hormones; deficiency → goiter; solved with iodized salt.

• Compounds

  • Chemical combination of ≥2 elements in fixed ratios (e.g., $H<em>2O$, $CH</em>4$, $H<em>2O</em>2$).

  • Properties differ radically from constituent elements (e.g., explosive $Na$ + toxic $Cl_2$ → table salt $NaCl$).

Atomic Structure & Isotopes

• Subatomic particles

  • Proton: positive, in nucleus, ~1 amu.

  • Neutron: neutral, in nucleus, ~1 amu (slightly heavier).

  • Electron: negative, orbitals/cloud, $\frac{1}{2000}$ amu.

• Atomic number $Z$ = # protons (unique identifier).

• Mass number $A$ = protons + neutrons (whole number).

• Isotopes: atoms of same element with different neutron counts → different $A$.

  • Heavy isotopes may be radioactive; nucleus decays, changing element.

  • $^{14}C$ → $^{14}N$ decay used in fossil dating & metabolic tracing.

  • Medical/biological tracers track metabolic pathways & diagnose disorders; radiation requires safety badges.

  • Half-life concept illustrated with $^{14}C$ limitations.

Electron Energy Levels & Orbitals

• Potential energy depends on distance from positive nucleus; farther electron = higher $PE$.

• Bohr model (simplified): shells

  • 1st shell: max 2 e⁻

  • 2nd shell: max 8 e⁻

  • 3rd shell: max 18 e⁻ (biology focuses on first 10)

  • 4th shell: max 32 e⁻

• Quantum/cloud model: orbitals are 3-D probability regions (spherical s, dumbbell p, etc.).

• Valence electrons (outer shell) dictate chemical reactivity; atoms seek noble-gas configuration (octet rule).

Chemical Bonds

Covalent Bonds (dominant in biomolecules)

• Formed by sharing valence electrons between non-metals.

• Single bond = 1 pair (e.g., $H<em>2$); double = 2 pairs (e.g., $O</em>2$); triple = 3 pairs (e.g., $N_2$).

• Structural formula uses dashes; molecular formula lists atom counts.

• Bonding capacity = number of unpaired valence electrons (H 1, O 2, N 3, C 4).

Electronegativity & Polarity

• Equal sharing → non-polar covalent (e.g., $CH_4$).

• Unequal sharing → polar covalent (e.g., $H_2O$); electrons pulled toward more electronegative atom (O, N, F, Cl).

• Generates partial charges (δ⁻ near EN atom, δ⁺ near H) producing molecular dipoles.

Ionic Bonds

• Complete electron transfer (metal → non-metal).

• Creates cations (+) and anions (–); strong electrostatic attraction forms salts/lattices (e.g., $Na^+Cl^-$ cubic crystal).

• In aqueous environments ionic bonds weaken; water’s polarity separates ions.

Weak Interactions (Critical in Biology)

• Hydrogen bonds: non-covalent attraction between $H$ covalently bound to $N/O/F$ and another $N/O/F$.

  • Represented with dotted lines; strongest of weak forces.

  • Crucial for water properties, DNA base pairing, protein folding.

• Van der Waals Interactions

  • London dispersion (non-polar) & dipole-dipole (polar) forces.

  • Gecko toe adhesion example.

• Cumulative weak forces stabilize 3-D shapes of proteins, nucleic acids, etc.

Molecular Shape & Biological Recognition

• Shape determined by bond geometry (e.g., linear $O<em>2$, bent $H</em>2O$ @ 104.5°, tetrahedral $CH_4$).

• Molecular mimicry: drugs emulate natural ligands; opiates (morphine, heroin) mimic endorphin binding motifs to alleviate pain.

Chemical Reactions & Equilibrium

• Reactants → Products; matter conserved via balancing coefficients.

• Photosynthesis: $6CO<em>2 + 6H</em>2O + \text{light} \rightarrow C<em>6H</em>{12}O<em>6 + 6O</em>2$ (builds molecules, stores energy).

• Cellular respiration = reverse (combustion, releases energy).

• Reversible reactions use equilibrium arrows; biological equilibria are dynamic, not static (cells die at static equilibrium).

• Rate influenced by reactant concentration.

Water: The Molecule of Life

Structure & Hydrogen Bonding

• Polar molecule; O δ⁻, two H δ⁺ → up to 4 hydrogen bonds per molecule.

• Comparison: $H<em>2S$, $H</em>2Se$ are gases at RT; water is liquid due to H-bond network.

Four Emergent Properties

1. Cohesive Behavior

   • Cohesion: H-bonding between water molecules; aids upward transport in plants.

   • Adhesion: H-bonding to other substances (xylem walls).

   • Surface tension: creates meniscus; supports water striders.

2. Temperature Moderation

   • High specific heat $c=1\,cal\,g^{-1}°C^{-1}=4.184\,J\,g^{-1}°C^{-1}$; buffers climate & organismal temps.

   • Heat absorbed to break H-bonds, released when bonds form.

   • High heat of vaporization (≈ 40.79 kJ g⁻¹); evaporative cooling → sweating & transpiration; steam burns severe.

3. Expansion Upon Freezing

   • Ice forms hexagonal lattice, lowers density → floats; insulates aquatic life at 4 °C; explains burst pipes in cold climates.

4. Versatility as a Solvent

   • Hydration shells form via ion-dipole attractions; dissolve salts, polar covalent molecules (glucose, ethanol), create aqueous solutions.

   • Hydrophilic vs hydrophobic domains; oil & water immiscible.

   • Colloids (blood, milk, smog) scatter light; stable dispersions.

Solutions & Concentrations

• Solution = solute + solvent (water in bio contexts).

• Mole: $6.02\times10^{23}$ particles; enables mole-to-mole comparisons.

• Molarity $M = \frac{\text{moles solute}}{\text{liters solution}}$.

Acids, Bases & pH

• Water auto-ionization: $2H<em>2O \rightleftharpoons H</em>3O^+ + OH^-$

  • Ion-product constant $K_W = [H^+][OH^-] = 10^{-14}$ (25 °C).

• pH definition: $\text{pH} = -\log[H^+]$.

  • Neutral (distilled water): $[H^+]=[OH^-]=10^{-7}$ M → pH 7.00.

  • Acidic: [H^+]>[OH^-] → pH < 7.

  • Basic (alkaline): [H^+]<[OH^-] → pH > 7.

• Strong acids/bases dissociate completely (single arrow); weak ones establish equilibrium (double arrow).

• Biological systems tightly regulate pH; drastic shifts impair metabolism.

Practical/Real-World Connections & Ethical Notes

• Iodized salt addresses public-health goiter; iron supplements combat anemia.

• Radioisotope usage requires radiation safety (cancer risk).

• Climate buffering by oceans demonstrates high water specific heat; environmental policy relevance.

• Opioid design & addiction: molecular mimicry has medical & societal implications.

Study Tips & Connections

• Recall general chemistry (electron shells, electronegativity, molarity) for smooth progress in biomolecules.

• Understand weak vs strong bonds: impacts protein folding, DNA stability, water behavior.

• Practice pH calculations, molarity conversions, and balancing reactions.

• Relate physical properties (specific heat, density) to ecological phenomena (lake turnover, coastal climate).