EL

bio- chapter 2

Matter, Elements, and Atoms

  • Matter: anything with mass and space; can be solid, liquid, or gas.
  • Element: pure substance that cannot be broken down (92 natural elements).
  • Compound: a substance made of 2+ elements in a fixed ratio.
  • Major elements (99% of the body): O, C, H, N, Ca, P.
  • Subatomic particles:
    • Proton (+): defines the element (atomic number).
    • Neutron (0): contributes to mass.
    • Electron (−): forms a cloud, involved in bonding.
  • Isotopes: same element, different neutrons.
  • Stable isotopes: nuclei don’t change.

Chemistry and Life

  • Many biological problems are chemical in nature.
  • Understanding chemistry is essential to understanding life.
  • Emergent properties: compounds differ from their elements (e.g., Na + Cl → NaCl).

Elements Essential for Life

  • Humans need 25 elements; plants need 17.
  • Trace elements (<0.01%): Fe, I, Zn, etc. → small amounts but critical.

Atoms and Subatomic Particles

  • Atom = smallest unit of matter that keeps the properties of an element.
    • Proton (+) → defines element (atomic number).
    • Neutron (0) → contributes to mass.
    • Electron (−) → forms cloud, involved in bonding.

Isotopes & Radioactivity

  • Radioactive isotopes: nuclei decay, release energy.
  • Iodine (I): thyroid hormone. Deficiency → goiter, developmental problems. Solution: iodized salt.
  • Iron (Fe): oxygen transport, energy processing.
  • Fluoride (F): strengthens teeth. Added to water/toothpaste.
  • Uses in research/medicine: tracers, medical imaging, cancer treatment.
  • Dangers of radiation: uses in fossil dating, medical imaging, cancer treatment.

2.4 Radioactive Isotopes—Help or Harm

  • Cells cannot distinguish isotopes → radioactive forms are used like normal elements.
  • Tracers: follow chemical processes (e.g., C-14 used to study photosynthesis).
  • PET scans: radioactive tracers (like PIB) detect brain activity/diseases (e.g., Alzheimer’s).
  • High exposure damages molecules (especially DNA).
  • Nuclear accidents (Chernobyl, Fukushima) caused deaths, evacuations, and long-term cancer risks.
  • Radon gas: natural radioactive gas, 2nd leading cause of lung cancer in the U.S. → homes tested for safety.

6. Trace Elements and Human Health

  • Deficiency = common nutritional disorder.
  • Excess = toxic, organ damage.
  • Benefits: reduced cavities.
  • Controversy: public health vs. individual rights.

7. Science & Society

  • Public health decisions (like water fluoridation) balance scientific evidence and social concerns.
  • Importance: citizens must evaluate evidence critically.

2.5 Electrons and Chemical Properties

  • Electrons determine chemical behavior.
  • Only electrons in the outermost shell (valence shell) matter for bonding.
  • Incomplete shells → atoms tend to react to complete shells.
  • Full shells → atoms are stable/inert (e.g., helium, neon, argon).
  • Give up electrons; Accept electrons; Share electrons.
  • These interactions form chemical bonds.

2.6 Chemical Bonds

2.6.1 Ionic Bonds

  • One atom transfers electrons to another.
  • Results in opposite charges (ions) attracting.
  • Example: Na⁺ + Cl⁻ → NaCl (table salt).
  • Ion = atom/molecule with an electrical charge from electron gain/loss.
  • Compound is electrically neutral overall, but stable in fixed ratios (NaCl = 1:1).
  • Bond strength depends on environment:
    • Dry crystal → strong (hard to break).
    • In water → weak (ions separate, salt dissolves).
  • Many drugs are manufactured as salts → stable when dry, dissolve easily in water.

2.6.2 Covalent Bonds

  • Atoms share electrons to complete outer shells.
  • Valence = bonding capacity (# of bonds possible).
  • Example: Carbon has valence of 4 → bonds with 4 hydrogens to form methane (CH₄).
  • Types of Covalent Bonds:
    • Nonpolar covalent bond: electrons shared equally. → Ex: C–H bonds.
    • Polar covalent bond: electrons shared unequally (due to electronegativity differences). → Ex: H₂O → oxygen more electronegative, attracts electrons → O slightly negative, H slightly positive.

2.7 Ionic Bonds

  • (Covered under 2.6.1; see above for details on electron transfer, ion formation, environment-dependent strength, and examples like NaCl.)

2.8 Hydrogen Bonds

  • Covalent bonds = strong bonds inside molecules.
  • Hydrogen bonds = weak bonds between molecules (or within large molecules).
  • H atoms bonded to O by polar covalent bonds.
  • Water is a polar molecule → partial charges: Oxygen slightly negative, Hydrogen slightly positive.
  • Opposite charges attract → weak hydrogen bonds between water molecules.

Biological significance

  • Each H₂O can form up to 4 hydrogen bonds with neighbors.
  • Shape and function of proteins.
  • Hold the two strands of DNA together.
  • Involved in gene expression and protein translation.

2.9 Chemical Reactions

  • Chemical reactions = breaking and making bonds, rearranging matter.
  • Law of conservation of matter: matter is not created or destroyed, only rearranged.
    • 2H₂ + O₂ → 2H₂O
    • Bonds in H₂ and O₂ broken, new bonds formed in water.
  • Photosynthesis (key life reaction):
    • 6\ CO2 + 6\ H2O → C6H{12}O6 + 6\ O2
    • Uses sunlight → converts carbon dioxide + water into glucose + oxygen.
  • In cells: thousands of reactions occur in watery environments.

2.12 Ice Floats Because It Is Less Dense than Liquid Water

  • Water exists in three states: vapor, liquid, and solid (ice).
  • Unique property: Ice is less dense than liquid water.
  • Reason:
    • As water freezes, hydrogen bonds form stable, spacious 3D crystals.
    • This arrangement spreads molecules apart → lower density.
  • Consequences:
    • Ice floats on liquid water.
    • Prevents entire lakes and oceans from freezing solid.
    • Floating ice insulates water below → life (fish, aquatic organisms) can survive under ice.
  • Extra: Freezing water expands and can crack boulders (because of expansion of ice within rock cracks).

2.13 Water Is the Solvent of Life

  • Solution = uniform mixture of 2+ substances.
  • Solvent: dissolving agent (water).
  • Solute: substance dissolved (salt, sugar, etc.).
  • Aqueous solution: solution where water is the solvent.
  • Why water is a versatile solvent:
    • Polarity of water molecules allows them to surround & separate solutes.
    • Positive H ends attract negative ions (Cl⁻).
    • Negative O end attracts positive ions (Na⁺).
  • Examples:
    • Salt dissolves as ions are separated by water molecules.
    • Sugar (polar molecule) also dissolves by forming hydrogen bonds with water.
    • Even large molecules (proteins) dissolve if they have ionic or polar regions.
  • Biological significance:
    • Water dissolves solutes essential for life.
    • Blood and most body fluids are aqueous solutions.
    • Plant sap and seawater also rely on water’s solvent ability.

2.14 The Chemistry of Life is Sensitive to Acidic and Basic Conditions

  • Water ionization: a tiny fraction dissociates into H⁺ (hydrogen ions) and OH⁻ (hydroxide ions).
  • These ions are highly reactive, so changes affect proteins & biomolecules.
  • Acids & Bases:
    • Acid = donates H⁺ (e.g., HCl in stomach acid).
    • Base = reduces H⁺ concentration; may donate OH⁻ or accept H⁺ directly (e.g., NaOH).
  • pH Scale:
    • Ranges 0 (acidic) → 14 (basic); 7 = neutral.
    • Each step = 10× change in H⁺ concentration.
    • Pure water: pH 7.
    • Human blood: pH 7.4 (life-threatening if below 7.0 or above 7.8).
  • Buffers:
    • Substances that resist pH changes.
    • Function: accept excess H⁺ or donate H⁺ when depleted.
    • Crucial for keeping pH stable in cells & blood.

2.15 Scientists Study the Effects of Rising CO₂ on Coral Reef Ecosystems

  • Ocean Acidification:
    • Oceans absorb ~25% of fossil fuel CO₂.
    • Dissolved CO₂ + H₂O → carbonic acid → lowers pH.
    • Ocean pH: dropped 0.1 unit in 420,000 years; may drop another 0.3–0.5 by 2100.
  • Impact on Corals:
    • Corals build skeletons via calcification: \mathrm{Ca^{2+} + CO3^{2-} → CaCO3}
    • H⁺ from acidification binds CO₃²⁻ → makes HCO₃⁻.
    • Result: 40% fewer carbonate ions by 2100 → weaker coral skeletons.
  • Controlled Experiments (Fig. 2.15A):
    • Independent variable: carbonate ion concentration.
    • Dependent variable: calcification rate.
    • Finding: lower carbonate → slower coral growth.
  • Field Studies (Fig. 2.15B):
    • Volcanic CO₂ seeps ("champagne reefs") naturally lower pH.
    • pH 7.8 vs. 8.1 → reduced coral diversity, fewer juveniles, weaker coral structures.
    • Reef ecosystems lose biodiversity & resilience.
  • Conclusion:
    • Multiple studies confirm: rising CO₂ threatens coral reefs & marine biodiversity.

2.16 The Search for Extraterrestrial Life Centers on the Search for Water

  • Why water?
    • Water’s unique properties (solvent, temperature buffer, density, etc.) are essential for life.
    • Searching for life = searching for water.
  • Evidence of Water on Mars:
    • Ice caps at poles.
    • 2008 Phoenix lander: ice under surface.
    • 2012 Curiosity rover: soil with high water content.
    • 2013 Opportunity rover: clay minerals (neutral-pH water once existed).
    • 2011 Mars Orbiter: streaks suggest seasonal melting streams.
    • 2015 NASA discovery: waterlogged molecules confirmed → liquid water on surface.
  • Significance:
    • Strong evidence for past or present life on Mars.
    • Discovery of life beyond Earth → new perspective on evolution.