Chapter 2 Notes: The Chemical Basis of Life (Campbell Biology 11th Edition)

Elements, Atoms, and Compounds

  • Life is chemistry: organisms are composed of matter and chemical reactions underlie all biological processes. Water is central to life; most body reactions occur in a water-rich environment (70%–95% water in cells).
  • Matter and states: solid, liquid, gas.
  • An element is a substance that cannot be broken down by ordinary chemical means. There are 92 naturally occurring elements; some additional synthetic elements exist.
  • Each element has a symbol (often derived from English, Latin, or German names).
  • A compound is a substance consisting of two or more elements in a fixed ratio. Compounds are more common in nature than pure elements.
  • Emergent properties: compounds have characteristics not present in their constituent elements. For example, table salt (sodium chloride, NaCl) forms edible salt that is not a property of sodium metal or chlorine gas.
  • Common elements in living matter: about 25 elements are essential to life. The four elements O, C, H, and N account for roughly 96% of body weight in humans and many organisms. They are the main ingredients of biological molecules (proteins, sugars, fats).
  • Table 2.1 (essential elements by weight):
    • Oxygen O: ~65.0%
    • Carbon C: ~18.5%
    • Hydrogen H: ~9.5%
    • Nitrogen N: ~3.3%
    • Calcium Ca: ~1.5%
    • Phosphorus P: ~1.0%
    • Potassium K: ~0.4%
    • Sulfur S: ~0.3%
    • Sodium Na: ~0.2%
    • Chlorine Cl: ~0.2%
    • Magnesium Mg: ~0.1%
    • Trace elements (less than 0.01% of body weight): boron, chromium, cobalt, copper, fluorine, iodine, iron, manganese, molybdenum, selenium, silicon, tin, vanadium, zinc.
  • Trace elements: essential in minute quantities; some are common additives in food and water (e.g., iron in cereals; iodine in salt; fluoride for dental health).
  • Fortified foods and labeling: foods may be fortified with minerals (e.g., iron in cereals) and vitamins; nutrition labels may show % Daily Value for added nutrients.

Atoms, Isotopes, and Atomic Structure

  • An atom is the smallest unit of an element that retains its properties. Atoms are extremely small (millions would fit across a length that is a few centimeters).
  • Subatomic particles relevant here:
    • Protons: + charge
    • Neutrons: no charge
    • Electrons: - charge
  • The nucleus contains protons and neutrons; electrons form a surrounding cloud at various energy levels.
  • Atomic number: number of protons in an atom. Defines the element (e.g., Helium has atomic number 2; Carbon has atomic number 6).
  • Mass number: sum of protons and neutrons in the nucleus.
  • Mass units: the dalton (approximately the mass of a proton or neutron).
  • Isotopes: atoms of the same element with the same atomic number but different mass numbers due to different numbers of neutrons. Examples with carbon:
    • Carbon-12: 6 protons, 6 neutrons, 6 electrons
    • Carbon-13: 6 protons, 7 neutrons, 6 electrons
    • Carbon-14: 6 protons, 8 neutrons, 6 electrons
    • Stability: ^12C and ^13C are stable; ^14C is radioactive.
  • Ions: atoms that have gained or lost electrons, acquiring a net electrical charge (e.g., Na^+ after losing one electron; Cl^- after gaining one electron).
  • Atomic mass is approximately the mass number (sum of protons and neutrons) because electrons contribute negligibly to atomic mass.
  • Important isotopes in biology: radioactive isotopes can be used as tracers in research (e.g., ^14CO2 to study photosynthesis) and in medicine (PET imaging with radiolabeled glucose or other molecules).

Radioactive Isotopes

  • Uses as tracers: biologically indistinguishable from nonradioactive isotopes but detectable by instruments.
  • Basic research: track molecules through chemical changes in organisms.
  • Medical diagnosis and treatment: tag compounds to visualize or target specific tissues; PET imaging detects metabolism.
  • Dangers: radiation can damage molecules, especially DNA; uncontrolled exposure is harmful (historical examples include Chernobyl, and radon exposure risks).
  • Example applications:
    • Carbon-13 or carbon-14 tracing in metabolic pathways
    • Iodine-131 for thyroid cancer treatment
    • PIB with radioactive label used in PET to visualize beta-amyloid plaques in Alzheimer's research
  • Safety note: small, controlled doses are used in medical contexts; exposure risks vs diagnostic benefits must be balanced.

Chemical Bonds and the Distribution of Electrons

  • Why atoms bond: atoms with incomplete outer shells interact to complete their valence shells, often forming stable structures held together by chemical bonds.
  • Electron shells and orbitals:
    • Electron shells around the nucleus can hold 2 electrons (first shell) or up to 8 in subsequent shells.
    • Each orbital holds up to 2 electrons.
  • Valence (bonding capacity): the number of additional electrons an atom needs to fill its outer shell. For main biological elements, typical valences are:
    • Hydrogen: 1
    • Oxygen: 2
    • Nitrogen: 3
    • Carbon: 4
  • Covalent bonds: strongest type of bond; atoms share electron pairs to fill valence shells, forming molecules (e.g., H2, O2, H2O, CH4).
  • Representations of molecules (examples with hydrogen):
    • Molecular formula: H2 (two hydrogen atoms)
    • Electron distribution: shows shared electron pairs
    • Structural formula: H–H (single covalent bond)
    • Space-filling model: spheres representing atoms and their relative sizes, showing molecule shape
  • Electronegativity: an atom's tendency to attract shared electrons.:
    • In O2, H2, CH4, electrons are shared relatively equally; bonds are nonpolar covalent when electronegativity differences are small.
    • In H2O, oxygen is highly electronegative and pulls electrons toward itself, creating a polar covalent bond. Oxygen becomes partially negative; hydrogens are partially positive, giving water a polar nature and bent shape.
  • Polar covalent vs nonpolar covalent:
    • Polar covalent bonds: unequal sharing of electrons (e.g., H2O).
    • Nonpolar covalent bonds: equal sharing (e.g., H2, O2, CH4).
  • Ionic bonds: transfer of electrons creates ions, which attract to form ionic compounds (e.g., NaCl). In NaCl, Na transfers one electron to Cl, producing Na^+ and Cl^-; ionic bonds form a crystal lattice; in water, these bonds can dissociate as ions separate in solution.
  • Hydrophobic vs hydrophilic interactions are related to polarity and solubility, but those terms are more about solvation in water and are not central to the basic bond types.

Hydrogen Bonds and Weak Interactions

  • Hydrogen bonds are weaker-than-covalent bonds but crucial in biology.
  • Formed between a partially positive hydrogen atom attached to an electronegative atom (like O or N) and a nearby electronegative atom in another molecule.
  • Water is an excellent example: each water molecule can form up to four hydrogen bonds (two as donors via its hydrogens and two as acceptors via the oxygen lone pairs).
  • Role in biology: stabilize protein structures, hold the two DNA strands together, contribute to water’s many life-supporting properties.

Chemical Reactions: Making and Breaking Bonds

  • Chemical reactions rearrange matter by breaking existing bonds and forming new ones, but do not create or destroy atoms.
  • Reactants vs products: the starting materials vs the substances formed.
  • Example: photosynthesis in plants. Overall simplified equation:
    6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2.
  • In photosynthesis, light energy powers the conversion of CO2 and H2O into sugar (glucose) and O2 as a by-product; atoms are conserved, but their arrangement changes.
  • In cellular metabolism, thousands of reactions occur in aqueous environments inside cells.

Water’s Life-Supporting Properties

  • Water’s emergent properties stem from its polarity and hydrogen bonding:
    • Cohesion: water molecules stick to each other via hydrogen bonds, contributing to surface tension and the transport of water in plants.
    • Adhesion: water also sticks to other surfaces (e.g., plant cell walls), aiding transport in narrow vessels.
    • Surface tension: high surface tension allows small organisms like water striders to walk on water.
  • Temperature regulation:
    • Water resists temperature change due to hydrogen bonding; energy must break bonds to raise temperature, and forming bonds releases energy when cooling.
    • This moderates climate and stabilizes ocean and freshwater temperatures, helping life.
    • Evaporative cooling: as water evaporates, the most energetic molecules leave, cooling the surface.
  • Water density and phase behavior:
    • Ice is less dense than liquid water because of its lattice structure formed by hydrogen bonds, causing ice to float.
    • Floating ice insulates bodies of water, protecting aquatic life in winter.
  • Water as solvent (universal solvent of life):
    • Water dissolves many substances due to its polarity. It can dissolve ionic compounds (e.g., NaCl) and polar molecules (e.g., sugars).
    • Hydration shells around ions and polar solutes enable biochemical reactions and transport in organisms.
    • An aqueous solution is one where the solvent is water.
    • Surface of dissolving salt: water molecules surround and separate ions; Na+ and Cl− are stabilized by hydration.
  • Acids, bases, and pH:
    • In water, some molecules dissociate into H+ (protons) and OH− (hydroxide).
    • Acid: donates H+ to solution (e.g., HCl → H+ + Cl−).
    • Base: accepts H+ or releases OH− (e.g., NaOH → Na+ + OH−; OH− + H+ → H2O).
    • pH scale: 0 (very acidic) to 14 (very basic), with 7 being neutral.
    • pH units describe 10-fold changes in H+ concentration: e.g., a solution at pH 2 has 10x more H+ than at pH 3.
    • Buffers help resist changes in pH by absorbing or releasing H+ as needed.
  • Acid precipitation and ocean acidification:
    • Emissions of sulfur oxides (SOx) and nitrogen oxides (NOx) react with water in the atmosphere to form acids; precipitation with pH < 5.2 harms lakes, forests, and soil chemistry.
    • Burning fossil fuels releases CO2; about 25% of human-emitted CO2 is absorbed by the oceans, lowering seawater pH and reducing carbonate ion availability needed for calcifying organisms (corals, shellfish).
    • Ocean acidification threatens marine food webs and biodiversity.
  • Extraterrestrial life and water:
    • Water is a key criterion in the search for life beyond Earth because of its life-supporting properties.
    • NASA discoveries: evidence for past water on Mars (Spirit and Opportunity rovers; Phoenix lander).
    • The presence of water is a strong indicator for environments that could support life, though not proof of life itself.

Connecting Concepts and Practice

  • Concept maps help tie together elements, atoms, bonds, and water properties.
  • Review concepts include: chemical bonds, molecular Representations, polarity, acids/bases, pH, buffers, and water’s role in temperature regulation and solvation.
  • Practice questions (selected examples) from the chapter review:
    • 3. Changing the number of protons would change the element (atomic number).
    • 4. Your body contains the smallest amount of which element? (answer: typically nitrogen or hydrogen depending on phrasing; the text emphasizes trace vs major elements; see Table 2.1 for emphasis on major elements.)
    • 5. A solution at pH 16 vs pH 8: pH 16 is outside the standard scale; the reference expects pH 1–14; interpret as 10^2 difference on a log scale in exam context.
    • 6. Most of the unique properties of water result from the fact that water molecules are polar and form hydrogen bonds.
    • 7. A sulfur atom has 6 electrons in its outer shell. As a result, it forms covalent bonds with other atoms with valence 2, 4, or 6 depending on the partner to complete its outer shell. (Sulfur commonly forms two to six bonds in compounds.)
    • 8. A trace element is required in very small amounts.
    • 9. A can of cola is an aqueous solution where water is the solvent and carbon dioxide makes the solution acidic (pH < 7).
      1. Radioactive isotopes can be used in medical studies because their radioactivity allows tracking of the isotope’s location and quantity in the body.
  • True/false practice items address basic corrections (e.g., table salts and sugars are compounds; the smallest unit of an element is an atom; ice floats due to lower density than liquid water; etc.).

Key Equations and Concepts (LaTeX)

  • Photosynthesis (overall):
    6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2.
  • Hydrogen ion and hydroxide ion in water:
    \mathrm{H_2O \rightleftharpoons H^+ + OH^-}.
  • Carbonic acid formation from CO2 in water (ocean chemistry):
    CO2 + H2O \rightarrow H2CO3 \rightarrow H^+ + HCO_3^-.
  • pH and hydrogen ion concentration relationship:
    \mathrm{pH = -\log_{10}[H^+]}.
  • Water’s density behavior: ice is less dense than liquid water due to hydrogen-bonded lattice formation (ice floats).
  • Ionic bond formation (NaCl example): Na transfers 1 electron to Cl, producing Na^+ and Cl^− which attract to form NaCl.

Connections to Real World and Implications

  • The chemistry of life explains why water is indispensable for biology, climate regulation, nutrient transport, and energy processing.
  • Trace elements, fortification of foods, and fluoridation of water illustrate public health applications of chemistry in biology.
  • Ocean acidification linked to anthropogenic CO2 has broad ecological and economic consequences for marine ecosystems and food webs.
  • The search for extraterrestrial life prioritizes the presence of liquid water as a key environmental requirement.

Notes on Figures and Tables Mentioned (Contextual)

  • Figure references: 2.1 (emergent properties of sodium chloride); 2.3 (isotopes of carbon and carbon-12, -13, -14); 2.5 (electron shells); 2.6 (covalent bonds representations); 2.7 (ionic bond formation); 2.8 (hydrogen bonds between water molecules); 2.9 (chemical reaction schematic); 2.10–2.16 (properties of water and environmental implications).
  • Table 2.1 (Elements in the human body by weight); Table 2.3 (isotopes of carbon); Table 2.6 (representations of common molecules).

Quick Reference: Terminology

  • Element: substance that cannot be broken down by ordinary chemical means.
  • Compound: substance composed of two or more elements in a fixed ratio.
  • Atomic number: number of protons; determines the identity of the element.
  • Mass number: total number of protons and neutrons.
  • Isotope: variants of an element with the same number of protons but different numbers of neutrons.
  • Ion: atom or molecule with a net electric charge due to gained or lost electrons.
  • Covalent bond: sharing of electron pairs between atoms.
  • Polar covalent bond: unequal sharing of electrons, resulting in partial charges.
  • Nonpolar covalent bond: equal sharing of electrons.
  • Ionic bond: transfer of electrons creating oppositely charged ions.
  • Hydrogen bond: weak bond between a hydrogen atom and an electronegative atom (e.g., O or N).
  • Acid: substance that donates H+ in solution.
  • Base: substance that accepts H+ or donates OH−.
  • pH: measure of how acidic or basic a solution is.
  • Buffer: substance that minimizes changes in pH.
  • Solvent: substance in which solutes are dissolved (water in biological systems).
  • Solute: substance dissolved in a solvent.
  • Aqueous solution: solution in which water is the solvent.
  • Photosynthesis: energy- and biomass-producing process in plants converting CO2 and H2O to glucose and O2.