Molecules of Life Powerpoint

Page 1: Molecules of life

  • Theme: atoms and electron arrangement underpin biology.
  • Electron shells (periodic system concept shown visually):
    • 1st shell: max 2 electrons. Example elements filling the 1st shell include Hydrogen (H, Z=1) and Helium (He, Z=2).
    • 2nd shell: fills with elements from Lithium (Li, Z=3) to Neon (Ne, Z=10).
    • 3rd shell: fills with elements from Sodium (Na, Z=11) to Argon (Ar, Z=18).
  • Note on the slide text quirks: it lists the sequence of elements and their shell occupation (e.g., 2 Li, 3 Be, 4 B, 5 C, 6 N, 7 O, 8 F, 9 Ne; 10 Na, 11 Mg, 12 Al, 13 Si, 14 P, 16 S, 17 Cl, 18 Ar) to illustrate how electrons populate shells as atomic number increases.
  • Key takeaway: shell capacity and order illustrate why atoms behave chemically the way they do; the pattern used in these slides is a simplified model for learning.
  • Summary concept: the arrangement of electrons in shells determines an element’s chemical properties and reactivity.

Page 2: The Biosphere and biological hierarchy

  • Hierarchy (from smallest to largest):
    • 1 Molecules (chlorophyll)
    • 2 Organelles (chloroplast)
    • 3 Cells
    • 4 Tissues
    • 5 Organs (leaf)
    • 6 Organisms (lupine)
    • 7 Populations (lupines)
    • 8 Communities
    • 9 Ecosystems (mountain meadow)
  • Emphasis: chemistry underpins all cellular processes; everything cells are and do has a molecular/chemical basis.

Page 3: States of matter in biology

  • All living organisms are made of matter that exists in three phases:
    • Solid
    • Liquid
    • Gas
  • Conceptual note: matter and its phases affect everything from cell structure to physiology.

Page 4: Composition of the human body by elements

  • Major elemental composition by mass:
    • Oxygen (O): 65\%
    • 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: < 0.01\% each (examples include Iron).
  • Overall message: 99% of the body is made of a small set of elements; most life is organized into compounds using these elements.
  • Note: Humans require 25 elements to live; elements are usually combined into compounds.

Page 5: Basic chemistry concepts

  • Opposite charges attract.
  • Subatomic particles:
    • Protons: + charge
    • Electrons: - charge
    • Neutrons: neutral
  • Definitions:
    • Element: a substance that cannot be broken down by chemical means (e.g., C, O, N).
    • Atom: the basic unit of an element.
  • Key relationships:
    • Atomic number Z = number of protons (determines element type).
    • Mass number A = Z + N, where N is the number of neutrons.
  • Connection between elements and subatomic particles underpins chemical behavior.

Page 6: Isotopes and radioactivity

  • Isotopes: same atomic number (same number of protons) but different mass number (different number of neutrons).
    • Examples: ^{12}C, \ ^{13}C, \ ^{14}C}
  • Some isotopes are unstable and radioactive; nucleus decays over time (e.g., ^{14}C).
  • Radioisotopes are valuable as biological tools (e.g., tracing molecules, dating, imaging).

Page 7: Electrons and chemical properties (initial shell concepts)

  • Number and location of electrons determine chemical properties.
  • 1st shell orbital concept; electrons occupy orbitals within shells.
  • Carbon example: number of electrons equals number of protons (in neutral atoms).
  • Electrons reside in orbitals; a maximum of 2 electrons can occupy the same orbital (Pauli exclusion principle in simplified terms for the first shell).
  • There are fixed numbers of orbitals in each energy level (shell).

Page 8: Valence electrons and shell descriptions

  • Outer (valence) electrons have the highest energy and largely determine reactivity.
  • Maximum valence electrons by shell (in the simplified model used here):
    • 1st shell: 2 e−
    • 2nd shell: 8 e−
    • 3rd shell: 8 e−
  • Atoms with fewer than the maximum valence electrons are typically reactive.
  • Example: Neon (Ne) has two filled shells (10 electrons total): 1st shell = 2 e−, 2nd shell = 8 e−.
  • Reiteration: electron distribution in shells/orbitals sets chemical properties.

Page 9: Chemical bonds and models

  • Bond: an attraction that holds atoms, ions, or molecules together; bonds shape molecules and store energy.
  • Common representations/models:
    • Space-filling model
    • Ball-and-stick model
    • Hybrid-orbital model (often shown with ball-and-stick superimposed)
  • Examples:
    • Water: H–O–H (H2O)
    • Methane: CH4
  • Key ideas:
    • Lone (unbonded) electron pairs influence geometry and bonding.
    • Bonding changes molecular shape and properties.

Page 10: Chemical reactions and bonds

  • Chemical reactions involve breaking and forming bonds.
  • Law of conservation of mass: matter cannot be created or destroyed; number and type of atoms balance between reactants and products.
  • Reactions rearrange matter by changing bonds and the positions of atoms, not by creating/destroying atoms.

Page 11: Water structure and Lewis representations

  • Hydrogen (H) has 1 valence electron; Oxygen (O) has 6 valence electrons.
  • First shell capacity: 2 electrons.
  • Approach to depict water:
    • Name and molecular formula: H2O
    • Space-filling model, electron distribution diagram, Lewis dot structure, and structural formula.
    • Water geometry: H–O–H with two lone pairs on O.

Page 12: Chemical equilibrium

  • Reversible reactions can reach dynamic equilibrium: no net change in concentrations; the system maintains a ratio at equilibrium.
  • Factors that influence equilibrium: temperature (thermal energy), concentration, and how favorable the reaction is.
  • Example: \ce{3 H2 + N2

Page 13: Covalent bonds

  • Covalent bonds form when atoms share electrons; outer orbitals merge to form molecular orbitals.
  • Strong bonds result from the mutual attraction between protons and shared electrons.
  • Biology-focused roles of covalent bonding: structure, storage, and the ability to form multiple bonds.
  • Multiple bonds: sharing more than one pair of electrons can form double or triple bonds.

Page 14–15: Electronegativity and bond polarity

  • Electronegativity influences covalent bond character.
  • Nonpolar covalent bonds: similar electronegativity; electrons shared equally; no permanent partial charges.
  • Polar covalent bonds: differences in electronegativity lead to unequal electron sharing; electrons spend more time around the more electronegative atom, creating partial charges (δ+ and δ−).
  • Visual cue: δ− on the more electronegative atom, δ+ on the less electronegative atom in a polar bond.

Page 16: Hydrogen bonds

  • Hydrogen bond: a weak, non-covalent interaction that can break and reform easily.
  • Requirements: a hydrogen atom bonded to a highly electronegative atom (such as O, N, or F) and a separate electronegative atom with a partial negative charge.
  • Examples: water (H2O) and ammonia (NH3) exhibit hydrogen bonding patterns (δ+ on H, δ− on the electronegative atom).

Page 17: Ionic bonds

  • Ionic bonds form when electronegativity differences are large enough to transfer electrons, creating ions.
  • Oppositely charged ions attract each other to form ionic bonds.

Page 18: Ionic compounds and biology

  • Example: Na+ and Cl− form ionic bonds in sodium chloride (NaCl).
  • Characteristics: ionic compounds are not discrete molecules with covalent bonds; they form strong lattice structures.
  • Biological relevance: signaling molecules, osmoregulation, structural components, and processes involving electron transfer.

Page 19: Water as a key biological solvent

  • All life depends on water.
  • Water content in humans: approximately 55%–60% of body mass (variations by sex and individual).
    • 55% for adult female, around 60% for adult male (typical ranges shown).

Page 20: Hydrogen bonding in water

  • Water forms hydrogen bonds via polar O–H bonds.
  • Each water molecule can participate in up to ~4 hydrogen bonds (two as hydrogen donors, two as acceptors).
  • Consequences:
    • Cohesion: water molecules stick to each other.
    • Water has high surface tension due to organized hydrogen bonding at interfaces.

Page 21: Surface tension and cohesion in water

  • High surface tension arises from the cohesive forces between water molecules at the air-water interface, pulling the surface into a tight film.
  • Result: water forms droplets and curves when in a cup, and water’s surface acts as a semi-rigid membrane.

Page 22: Adhesion and interactions with surfaces

  • Water adheres to other surfaces (adhesion), enabling wetting of materials like cellulose.
  • Example shown: hydrogen-bonding network between water and cellulose surfaces.

Page 23: Water as solvent and hydration shells

  • Water as solvent for life: solvent, solute, and aqueous solutions.
  • Hydration shells: water molecules surround ions (e.g., Na+, Cl−) and small molecules, disrupting some ionic interactions to stabilize dissolved species.
  • Process: dissolution involves water interacting with solute via hydrogen bonding and polarity, without breaking all covalent bonds within solutes.

Page 24: Hydrophilic substances and dissolution

  • Hydrophilic substances interact favorably with water.
  • Polar molecules dissolve in water; covalent bonds within solutes do not necessarily break during dissolution (water stabilizes the separated ions/molecules through hydrogen bonding and dipole interactions).

Page 25: Solubility of large molecules

  • Large, hydrophilic molecules (e.g., certain proteins) can be brought into solution by water depending on their surface chemistry and charge distribution.
  • The slide imagery suggests proteins with regions that attract water (solvation).

Page 26: Hydrophobic compounds and amphipathic molecules

  • Hydrophobic compounds are nonpolar and resist dissolution in water (e.g., fats, oils, waxes).
  • Amphipathic molecules have both hydrophilic and hydrophobic parts and can align at interfaces (e.g., phospholipids in membranes).

Page 27: Ice and liquid water; density and habitat implications

  • Hydrogen bonds in ice are more stable than in liquid water, making ice less dense than liquid water.
  • Consequence: ice floats, insulating liquid water below and providing habitat for aquatic life in cold climates.

Page 28: Thermal energy and related concepts

  • Key terms:
    • Kinetic energy: energy of motion
    • Thermal energy: related to temperature and molecular motion
    • Temperature: measure of average kinetic energy
    • Heat: transfer of thermal energy between bodies
  • Units used: calories (cal) and kilocalories (kcal)

Page 29: Water’s high specific heat and thermal buffering

  • Specific heat of water: 1\ \text{cal}\, (\text{g} \cdot {}^\circ\mathrm{C})^{-1} (per gram per degree Celsius), which is high.
  • Implications:
    • Water resists quick temperature changes (thermodynamic buffer).
    • Moderates climate, stabilizes body temperature, and influences biochemical reaction rates.

Page 30: Water autoprotolysis and acid-base basics

  • Water can dissociate into ions: hydroxide OH− and hydrogen H+.
  • The chemistry of life is sensitive to acidic/basic conditions; shifting ion concentrations alters biochemical processes.
  • Relevant equilibrium: \mathrm{H_2O \rightleftharpoons H^+ + OH^-} and pH dependence.

Page 31: pH, hydrogen ion concentration, and water activity

  • Relationship between hydrogen ion concentration and pH:
    • pH = -\log[H^+]
  • Water’s autoionization product is constant: [H^+][OH^-] = 10^{-14} (at 25°C).
  • If [H^+] = 10^{-7}\,\text{M}, then pH = 7 (neutral).
  • As a logarithmic scale, each unit change in pH represents a 10-fold change in the hydrogen ion concentration.
  • Example points on the scale: neutral around 7; acidic below 7; basic above 7.

Page 32: A representative pH scale with common substances

  • The scale is shown from highly basic to highly acidic:

    • 14: Very basic (e.g., strong bases like concentrated NaOH)
    • 13: Household bleach (highly basic)
    • 12–11: Very basic to strongly basic (amines like ammonia fall here)
    • 10–9: Basic solutions (bleach remnants, some cleaners)
    • 9–8: Seawater to mild bases (baking soda around 9; Seawater around ~8)
    • 7: Neutral (pure water; human blood near 7.35–7.45)
    • 6–5: Fresh water toward mildly acidic conditions
    • 4–3: Apple juice to stomach fluids (acidic range)
    • 0–2: Battery acid to strong acids (e.g., hydrochloric acid in gastric juice ~1–2)

  • Practical takeaway: Small changes in pH reflect large shifts in hydrogen ion concentration and can dramatically affect biochemical processes.

  • Overall integration: The slides collectively establish how basic chemical principles—atomic structure, bonding, water chemistry, and pH—underlie biology, from molecular interactions to organismal physiology and ecosystems.