GP

Chemisty of life

The Four Postulates of Atomic Theory

  • All matter is composed of atoms.
  • Atoms of a given element differ from atoms of all other elements.
  • Chemical compounds consist of atoms combined in specific ratios.
  • Chemical reactions change only the way the atoms are combined in compounds; the atoms themselves are unchanged.

Major Elements of the Body

  • Major elements by percent of total body weight (approximate):
    • Oxygen: 65.0\%
    • Carbon: 18.6\%
    • Hydrogen: 9.7\%
    • Nitrogen: 3.2\%
    • Calcium: 1.8\%
    • Phosphorus: 1.0\%
    • Potassium: 0.4\%
    • Sodium: 0.21\%
    • Chlorine: 0.21\%
    • Magnesium: 0.11\%
    • Sulfur: 0.05\%
    • Iron: 0.03\%
    • Iodine: 0.03\%
  • These four elements (O, C, H, N) make up about 96.5\% of body weight.

Atoms and Subatomic Particles

  • Atoms are building blocks of elements; smallest unit that retains chemical properties.
  • Atoms consist of two regions:
    • Nucleus (center) contains protons and neutrons (collectively, nucleons).
    • Electrons occupy orbit around the nucleus in electron clouds/orbitals.
  • Subatomic particles:
    • Protons (positive charge) and neutrons (neutral) in the nucleus.
    • Electrons (negative charge) in orbit around the nucleus.
  • Mass units: protons and neutrons are roughly 1 atomic mass unit (amu); electrons are ~1836 times lighter than protons/neutrons.
  • Protons and electrons have equal magnitude charges but opposite signs; atoms are electrically neutral overall.

Atomic Structure and Isotopes

  • Atomic Number (Z): number of protons in the nucleus; defines the element.
  • Mass Number (A): total number of protons and neutrons in an atom.
  • Isotopes: atoms of the same element with different numbers of neutrons; have same Z but different A.
  • Radioactive isotopes decay spontaneously, releasing particles and energy; applications include dating fossils, tracing metabolic processes, and diagnostic medical imaging.
  • Example: Chlorine isotopes and carbon dating concept discussed.

Isotopes and Atomic Mass (Weighted Averages)

  • Atomic weight is the weighted average mass of an element’s atoms, accounting for the natural abundances of its isotopes:
    • Formula: ext{Atomic weight} = \sumi (fi \times mi) where fi is the fractional abundance and m_i is the mass of isotope i.
  • Example: Chlorine exists as two isotopes with abundances and masses:
    • ^{35}\text{Cl}: abundance 75.77\%, mass 34.97\text{ amu}
    • ^{37}\text{Cl}: abundance 24.23\%, mass 36.97\text{ amu}
    • Atomic weight ≈ (0.7577)(34.97) + (0.2423)(36.97) = 35.45\text{ amu}
  • Atomic weight requires knowledge of each naturally occurring isotope’s mass and percent abundance.

The Periodic Table and Electron Configuration

  • Periodic table organizes elements in 7 periods (rows) and 18 groups (columns); groups share similar valence electron counts.
  • Atomic mass (approximate) appears below the element symbol; atomic number appears above.
  • Representative element examples and masses (selected):
    • Hydrogen (H): atomic mass ~ 1.00794\text{ amu}
    • Helium (He): ~ 4.00260\text{ amu}
    • Lithium (Li): ~ 6.941\text{ amu}
    • Beryllium (Be): ~ 9.01218\text{ amu}
    • Carbon (C): ~ 12.0107\text{ amu}
    • Nitrogen (N): ~ 14.0067\text{ amu}
    • Oxygen (O): ~ 15.9994\text{ amu}
    • Neon (Ne): ~ 20.1797\text{ amu}
  • The periodic table shows both atomic number Z and approximate atomic mass for each element.

Atomic Structure: Shells, Subshells, and Orbitals

  • Electronic structure is organized into shells, subshells, and orbitals:
    • Shells: numbered 1, 2, 3, …; farther shells are larger and can hold more electrons.
    • Subshells: s, p, d, f in order of increasing energy within a shell.
    • Orbitals: regions where electrons are most likely to be found; each orbital holds up to 2 electrons.
  • Maximum electrons per shell (capacity):
    • General rule: the shell capacity is 2n^2\$, where n is the shell number.
    • Examples: 1st shell: 2 electrons; 2nd shell: 8 electrons; 3rd shell: 18 electrons; 4th shell: 32 electrons.
  • Subshell capacities:
    • s subshell: 1 orbital ⇒ 2 electrons
    • p subshell: 3 orbitals ⇒ 6 electrons
    • d subshell: 5 orbitals ⇒ 10 electrons
    • f subshell: 7 orbitals ⇒ 14 electrons
  • The arrangement of electrons in shells and subshells is described by the electron configuration rules (Aufbau principle, Hund’s rule, Pauli exclusion).
  • Valence shell: outermost, highest-energy shell; contains valence electrons that largely determine chemical properties.
  • Periodic table organization relates to valence electrons; elements in the same group have the same number of valence electrons.

Electron Configuration and Rules

  • Electron configuration describes the exact arrangement of electrons in shells and subshells.
  • Rules for filling orbitals:
    • Electrons occupy the lowest-energy orbitals available first (Aufbau principle).
    • Each orbital holds up to two electrons with opposite spins (Pauli exclusion).
    • If two or more orbitals have the same energy, electrons occupy them singly first (Hund’s rule).
  • Order of orbital energy levels follows a specific sequence; a mnemonic is sometimes used to remember the order (e.g., 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p).

Electron Distribution and the Octet Rule

  • Valence electrons in the outer shell tend to achieve a noble gas configuration (octet, or 2 electrons for hydrogen) via covalent bonding or ionic bonding.
  • The octet rule guides formation of many common compounds, especially among main-group elements.
  • For many elements, achieving a full outer shell leads to high stability; exceptions exist (e.g., some compounds involving d-orbitals).

Covalent Bonds, Molecules, and Bonding Concepts

  • Covalent bond: achieved when atoms share electrons; forms molecules.
  • A molecule is a group of atoms held together by covalent bonds.
  • Main-group elements tend to achieve an octet (or 2 for H) via covalent bonds to fill valence shells.
  • Bond order and bond types:
    • Single bond: one shared electron pair.
    • Double bond: two shared electron pairs.
    • Triple bond: three shared electron pairs.
    • Represented in drawings by 1, 2, or 3 lines between atoms.
  • Polar covalent bonds vs nonpolar covalent bonds:
    • If atoms are identical, electrons are shared equally (nonpolar).
    • If atoms are different, electrons are shared unequally (polar covalent) due to electronegativity differences.
  • Electronegativity: ability of an atom to attract electrons in a bond.
    • Fluorine is most electronegative (value typically set to 4 in many scales).
    • Top-right elements are generally more electronegative; noble gases are not assigned values.
  • Bond polarity and dipoles:
    • Polar covalent bonds create partial charges on atoms (δ+ and δ−).
    • A molecule can be polar if the sum of bond dipoles and lone-pair contributions does not cancel out.
    • Symmetrical molecules with polar bonds may be nonpolar overall because dipoles cancel (e.g., CO₂, CCl₄).
  • Example diatomic molecules:
    • H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂ (color, reactivity variations noted).
  • Electronegativity differences and bond character:
    • ΔEN < 0.5 => nonpolar covalent bonds.
    • 0.5 ≤ ΔEN < 1.9 => polar covalent bonds.
    • ΔEN ≥ 2 => ionic bonds.
    • There is no sharp boundary; many bonds are intermediate in character.

Ions and Ion Formation

  • Ions are charged particles formed by gaining or losing electrons.
  • Cation: positively charged ion formed by loss of electrons (e.g., Na⁺).
  • Anion: negatively charged ion formed by gain of electrons (e.g., Cl⁻).
  • Alkali metals (group 1) have low ionization energies and readily form cations by losing one electron.
  • Halogens (group 17) have high electron affinities and readily form anions by gaining an electron.

Ionic Bonds and Crystal Lattices

  • Ionic bond: electrostatic attraction between ions of opposite charge in a crystal structure.
  • Ionic compounds: substances held together by ionic bonds (e.g., NaCl).
  • In NaCl, each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺ in a crystal lattice.

Acids, Bases, and Aqueous Chemistry

  • Acids donate protons (H⁺) in water; bases donate hydroxide ions (OH⁻) in water.
  • Examples:
    • HCl → H⁺ + Cl⁻
    • NaOH → Na⁺ + OH⁻
  • Common acids and their conjugate bases (ions):
    • Acetic acid: CH₃COOH; acetate: CH₃COO⁻
    • Carbonic acid: H₂CO₃; bicarbonate: HCO₃⁻; carbonate: CO₃²⁻
    • Hydrochloric acid: HCl; chloride: Cl⁻
    • Nitric acid: HNO₃; nitrate: NO₃⁻
    • Nitrous acid: HNO₂; nitrite: NO₂⁻
    • Phosphoric acid: H₃PO₄; dihydrogen phosphate: H₂PO₄⁻; hydrogen phosphate: HPO₄²⁻; phosphate: PO₄³⁻
    • Sulfuric acid: H₂SO₄; hydrogen sulfate: HSO₄⁻; sulfate: SO₄²⁻
  • Water autoionization: ext{H}2 ext{O} + ext{H}2 ext{O}
    ightleftharpoons ext{H}_