General Bio - The Chemical Foundation of Life (Ch. 2)

Page 1

  • Life is composed of matter; matter is anything that occupies space and has weight (e.g., air, living organisms).
  • Elements are unique forms of matter; have chemical and physical properties; each element is designated by a chemical symbol (e.g., S for Sulfur, Ca for Calcium). Common living-element symbols: Carbon (C), Oxygen (O), Hydrogen (H).
  • Atoms are the building blocks of elements. An atom is the smallest unit of matter that retains all chemical properties of an element.
  • An atom contains two regions:
    • Nucleus: center of the atom that contains protons and neutrons.
    • Orbitals (electron shells) surrounding the nucleus that hold electrons.
  • The symbol e⁻ is used to denote an electron.
  • Subatomic particles: protons, neutrons, and electrons.
  • Atomic Mass Unit (amu) = 1 dalton; defined as the mass of a proton. The mass of an electron is negligible in most calculations (often approximated as zero relative to protons/neutrons).
  • An approximate relationship: one neutron weighs about 1836 times as much as one electron. In other words, mn ≈ 1836 me. [Transcript notes give ~1800; standard value ≈1836].
  • An element is defined by the number of protons it has; changing the number of protons changes the element (i.e., the element’s identity).
  • Helium example: Helium has 2 protons, 2 neutrons, and 2 electrons; overall charge is neutral.
  • Protons and neutrons reside in the nucleus; electrons occupy orbitals around the nucleus.
  • Sub-atomic particles have distinct properties and masses; nuclear mass is largely concentrated in the nucleus; electrons contribute negligibly to atomic mass.
  • Quick recap of key symbols: p⁺ (protons), n⁰ (neutrons), e⁻ (electrons); atomic number Z = number of protons; atomic mass number A ≈ Z + N (where N is number of neutrons).
  • Quick reminder: AMU is used to express atomic/molecular masses, not the number of atoms or molecules themselves.
  • Connections to broader context: understanding atomic structure underpins chemical bonding, isotopes, isotopic abundance, and techniques like radiometric dating.
  • Quick conceptual link to real-world relevance: isotopes (stable vs. radioactive) are used in dating fossils, tracing biological processes, and understanding environmental changes.
  • Note on potential transcript ambiguity: some phrasing (e.g., “e’” for electron, or “James Harris discovered Berklienium-Black man Electron Shells”) appears miswritten; core concepts align with standard atomic theory (Bohr model, electron shells). The corrected mainstream interpretation is provided in later notes where appropriate.

1 amu=1 dalton1\ \mathrm{amu} = 1\ \mathrm{dalton}

  • The overall mass contribution of electrons is tiny compared to protons and neutrons, so atomic mass primarily reflects protons + neutrons.
  • Atomic mass unit (amu) is used to express both atomic and molecular masses; molecular mass in amu corresponds numerically to molar mass in g/mol for a given substance.

Page 2

  • Atomic number (Z): the number of protons in an atom; each element has a distinct atomic number.

  • The arrangement and identity of subatomic particles determine the element’s properties and behavior in reactions.

  • Atomic mass number (A): the total number of protons + neutrons in the nucleus; approximately equal to Z + N.

  • Neutrons (N) can vary within an element, producing isotopes.

  • Isotopes: forms of the same element (same Z) with different numbers of neutrons (different mass numbers A).

  • For many calculations (especially using a rough mass), electrons are not included in the mass calculation; mass is largely due to protons and neutrons.

  • Atomic mass is typically expressed in amu; isotopes differ in neutron count, hence mass number.

  • The carbon-centered statement: all life on Earth is organized around carbon; not all carbon-containing substances are organic (e.g., diamonds are pure carbon but can be considered inorganic in some contexts).

  • Neutron number can be determined by N = A - Z.

  • Example concept: If an element has Z = 6 (carbon), its isotopes differ in N (neutrons).

N=A−Z(neutrons)N = A - Z\quad\text{(neutrons)}

  • Note on carbon as life’s backbone: carbon’s versatile bonding underpins organic chemistry and biochemistry.

Page 3

  • Carbon (Z = 6) has two stable isotopes with mass numbers 12 and 13.
    • 12C: mass number A = 12; 12C has 6 neutrons (A - Z = 12 - 6 = 6).
    • 13C: mass number A = 13; 13C has 7 neutrons (A - Z = 13 - 6 = 7).
  • Natural abundance of carbon isotopes in nature approximates: ~99% 12C, ~1% 13C, with trace amounts of other (unstable) isotopes like 14C.
  • 12C atomic mass is 12 amu; however, the overall atomic mass of carbon (including isotopes) is about 12.011 amu12.011\ \mathrm{amu} due to the presence of 13C and 14C isotopes.
  • Isotopes: forms of an element with different numbers of neutrons, hence different mass numbers.
  • Hydrogen isotopes (all with Z = 1):
    • 1H: A = 1, N = 0
    • 2H (deuterium): A = 2, N = 1
    • 3H (tritium): A = 3, N = 2
  • Although all three isotopes have the same Z = 1, their masses differ, giving different atomic masses.
  • Because natural hydrogen includes all three isotopes, the average atomic mass of hydrogen is about 1.008 amu1.008\ \mathrm{amu}.
  • Radioisotopes: unstable isotopes that emit neutrons, protons, and electrons (radiation).
  • Radiometric dating leverages radioactive decay to estimate age.
    • Example: 14C decays to 14N over time (beta decay).
    • Researchers compare atmospheric 14C with fossil 14C to estimate fossil ages.
    • Carbon dating is effective for remains younger than ≈ 50,000 years.
    • For older objects, dating relies on isotopes of uranium or other elements.
    • Practical implication: carbon dating helps reconstruct past environments and biological timelines.

Page 4

  • The Periodic Table organizes elements; vertical columns (groups) have similar properties.
  • Group 18: Noble gases; characterized by a full valence shell; generally inert due to complete electron shells.
  • Historical note (transcript): “James Harris discovered Electron Shells and The Bohr Model” and a mention of Berklienium (likely a misstatement). In established history, Bohr proposed the Bohr model; electron shells are a core concept in atomic structure.
  • Core ideas (as intended in transcript):
    • The nucleus contains protons (p⁺) and neutrons (n⁰); electrons (e⁻) reside in electron shells around the nucleus.
    • In a neutral atom, the number of protons equals the number of electrons (Z = number of e⁻).
    • The Bohr model posits protons in the nucleus and electrons in circular orbits at fixed distances from the nucleus; this simplified picture introduces the concept of energy levels (shells).
    • Real atoms have more complex electron distributions (orbitals) beyond simple circular orbits.
  • Electron shells (energy levels): electrons occupy shells in order of increasing energy, filling the closest shell first before moving outward.
  • Each shell has a capacity: the first shell can hold up to 2 electrons; the second up to 8; subsequent shells hold more according to quantum rules (not all details in transcript, but implied by the shell capacity idea).

Page 5

  • Bohr model details and limitations:
    • Bohr model is an early, simplified depiction; it is incomplete because electrons do not move in fixed planet-like orbits in reality.
    • Electron orbitals are complex shapes that describe probability distributions of where electrons are likely to be found around the nucleus.
  • Electron shells and orbitals:
    • Electron shells (energy levels) are discrete distances from the nucleus.
    • Orbitals describe spatial distribution of electrons within those energy levels.
  • Neutral atoms: number of protons equals number of electrons (Z = number of e⁻).
  • Electrons determine how atoms interact; chemical reactions are changes in the distribution of electrons between atoms, leading to rearrangement of atoms in molecules.
  • Definitions:
    • Reactants: substances used at the beginning of a reaction.
    • Products: substances formed at the end of a reaction.
  • Bonding:
    • Chemical bond: an attractive force that links atoms together to form molecules and compounds.

Page 6

  • Electron filling and valence:
    • Outer shell (valence shell) determines chemical properties; stability is greatest when the valence shell is filled.
    • Group 18 noble gases possess full valence shells and are particularly stable.
  • Shell capacities (simplified, per transcript):
    • 1st shell: up to 2 electrons.
    • 2nd shell: up to 8 electrons.
    • Outer shell stability arises when these capacities are met.
  • Ionic bonding example:
    • Sodium (Na) and Chloride (Cl) atoms can transfer/share electrons to achieve filled outer shells, producing table salt (NaCl).

Page 7

  • Summary of Bohr model limitations:
    • Bohr models are incomplete because electrons are not in fixed, planet-like orbits.
    • Modern description uses electron orbitals (complex shapes) that describe spatial distributions.
  • Electron behavior governs chemical interactions:
    • Chemical reactions involve redistribution of electrons between atoms and consequent rearrangement of atoms in molecules.
  • Key terms:
    • Chemical bond: an attractive force that links atoms together to form molecules and compounds.
    • Reactants: substances present at the start of a reaction.
    • Products: substances formed at the end of a reaction.

Page 8

  • Chemical formulas vs structural formulas:
    • Chemical formula: indicates the number of each type of atom in a molecule.
    • Structural formula: illustrates how atoms are arranged in the molecule.
  • Mole (mol): a unit of amount of substance.
  • Stoichiometry basics:
    • Example equation: for hydrogen peroxide decomposition, 2 H₂O₂ → 2 H₂O + O₂ (balanced equation).
    • In real scenarios, reactions involve many millions of molecules; lab practice uses mole quantities to scale.
    • When starting with 2 moles of H₂O₂, products are 2 moles of H₂O and 1 mole of O₂.
  • Conceptual note: Equation can be viewed as representing amount of a particular element/molecule in a reaction, not a single molecule.

Page 9

  • Mole concept (mole):
    • A mole is defined as containing exactly 6.02 × 10²³ entities (atoms or molecules).
    • Avogadro's number: NA=6.02×1023N_A = 6.02\times 10^{23}.
  • Lab note (FYI): 1 mol of a substance has a mass in grams equal to its molecular mass in amu.
  • Example: Hydrogen peroxide (H₂O₂) has a molecular mass of approximately M(H<em>2O</em>2)=34 amuM(\mathrm{H}<em>2\mathrm{O}</em>2) = 34\ \mathrm{amu}
    • Therefore, 1 mol of H₂O₂ has a mass of about 34  g34\ \,\mathrm{g}.
  • Practical lab calculation:
    • Weigh 68 g of H₂O₂ (which is 2 mol) and you should obtain 2 mol of H₂O (mass 36 g) and 1 mol of O₂ (mass 32 g).
  • Note on terminology:
    • Molecular mass is sometimes called “molecular weight”; conceptually, the mole relates mass in grams to molecular mass in amu.
  • Final takeaway:
    • A mole corresponds to the AMU of the chemical and provides a bridge between atomic-scale masses and laboratory-scale quantities.