Foundations of Chemistry – Comprehensive Study Notes

Branches of Chemistry

  • FIVE principal branches → Organic, Inorganic, Analytical, Physical, Biochemistry. Chemistry is called “the central science” because each overlaps with biology, physics, geology, medicine, materials science, etc.
  • ORGANIC CHEMISTRY
    • Studies structure, properties, synthesis of C–H compounds.
    • Key sub-fields: medicinal, organometallic, polymer, physical-organic, stereochemistry.
    • Examples: drug design, Grignard reagents, nylon, SN1/SN2 kinetics, chirality in thalidomide.
  • INORGANIC CHEMISTRY
    • Concerned with all non-organic substances: elements, minerals, catalysts, metals, crystals.
    • Sub-branches: bioinorganic (metallo-enzymes), geochemistry, nuclear, solid-state.
  • ANALYTICAL CHEMISTRY
    • Qualitative & quantitative determination of composition.
    • Applications: forensics, environmental monitoring, bioanalytical drug screens.
  • PHYSICAL CHEMISTRY
    • Relates molecular structure ↔ observable physical properties.
    • Sub-disciplines: photochemistry, surface chemistry, chemical kinetics, quantum chemistry, spectroscopy.
  • BIOCHEMISTRY
    • Chemical processes in living organisms.
    • Components: molecular biology, genetics, pharmacology, toxicology, clinical & agricultural biochemistry.

Phases & States of Matter

  • Matter: anything with mass & occupies space.
  • Atom = smallest unit retaining chemical identity; nucleus (p⁺, n⁰) + electron cloud (e⁻).
  • Natural states: Solid, Liquid, Gas, Plasma. Man-made: Bose–Einstein Condensate (BEC).
  • SOLID
    • Particles fixed, vibrate; definite shape & volume; high density; low KE.
  • LIQUID
    • Particles flow; definite volume, indefinite shape; nearly incompressible.
  • GAS
    • Large intermolecular distances; high KE; no fixed shape/volume; compressible.
  • PLASMA
    • Ionized gas; high KE; prevalent in stars & neon signs.
  • BEC
    • Achieved near 0\,\text K; atoms coalesce into one quantum state ("super-atom"); used to study quantum mechanics, simulate black-hole conditions.
  • Phase changes (6): freezing, melting, condensation, vaporization, sublimation, deposition.

Classification of Matter: Pure Substances vs Mixtures

  • PURE SUBSTANCE → constant composition.
    • Element (e.g., \mathrm{O_2}) or compound (e.g., \mathrm{NaCl}).
  • MIXTURE → two or more substances physically combined; separable.
    • Heterogeneous (visible phases, ex: veggie soup).
    • Homogeneous / solution (uniform, ex: salt water).
  • Key takeaways:
    • Pure substances: fixed properties.
    • Mixtures retain individual component properties.

Properties & Changes of Matter

  • Physical properties: observed without identity change (color, density, m.p.).
  • Chemical properties: describe reactivity (flammability, rusting).
  • Intensive vs Extensive:
    • Intensive independent of amount (density, T).
    • Extensive depend on amount (mass, volume).
  • Physical change: no composition change (ice → water).
  • Chemical change / reaction: new substances formed (H₂ + O₂ → H₂O).

Energy: Forms & Fundamental Laws

  • Energy: ability to do work or transfer heat.
  • Kinetic Ek (motion) vs Potential Ep (position).
  • Eight common forms:
    1. Thermal (heat)
    2. Chemical (bond energy)
    3. Nuclear (fission/fusion)
    4. Electrical (moving charges)
    5. Radiant (EM waves)
    6. Sound (vibrational waves in matter)
    7. Elastic (deformed objects)
    8. Gravitational (height).
  • Law of Conservation of Energy: \Delta E{system}+\Delta E{surroundings}=0.
  • Law of Definite Proportions & Multiple Proportions govern fixed mass ratios in compounds.
  • Energy diagrams: exothermic (releases \Delta H

Atomic Structure & Isotopes

  • Atomic number Z = #protons; Mass number A = p⁺ + n⁰.
  • Isotopes: same Z, different A (e.g., ^{12}\text C vs ^{14}\text C).
    • Example: ^{14}\text C has 6 p⁺, 8 n⁰.

Chemical Nomenclature (Key Rules)

  • Elements: symbols (Fe, Na from Latin natrium).
  • Cations
    • Metals keep name (Na⁺ sodium ion).
    • Variable charge: Stock method Fe²⁺ iron(II); Classical ferrous.
  • Anions
    • Monoatomic: end in –ide (Cl⁻ chloride).
    • Polyatomic oxyanions: –ate (more O), –ite (less O); per– (one more), hypo– (one less).
    • Add H⁺: hydrogen carbonate \mathrm{HCO_3^-}.
  • Ionic compounds: cation name + anion name (Cu(ClO₄)₂ copper(II) perchlorate).
  • Acids:
    • Anion –ide → hydro___ic acid (HCl).
    • Anion –ate → ___ic acid (HNO₃ nitric).
    • Anion –ite → ___ous acid (HNO₂ nitrous).

Measurements in Chemistry (SI Units & Prefixes)

  • Base SI units used in chemistry: meter m (length), kilogram kg (mass), second s (time), kelvin K (temperature), mole mol (amount).
  • Common prefixes: \text{k}=10^3, \text{m}=10^{-3}, \mu=10^{-6}, \text{n}=10^{-9}, \text{p}=10^{-12}.
  • Temperature conversions:
    • K = ^\circ C + 273.15.
    • ^\circ C = \dfrac{(^\circ F-32)\,5}{9}.
  • Mass vs Weight: mass (kg) invariant; weight = force F=mg (newton).

Volume, Density & Significant Figures

  • Volume units: 1\,L = 1000\,cm^3 = 10^{-3}\,m^3; 1\,mL = 1\,cm^3.
  • Density \rho = \dfrac{m}{V}, typical unit g\,cm^{-3}.
  • Significant-figure rules:
    1. Non-zero digits significant.
    2. Captive zeros significant.
    3. Leading zeros not significant.
    4. Trailing zeros significant only if decimal shown.
  • Calculations:
    • Add/Subtract → limit to least precise decimal place.
    • Multiply/Divide → limit to least number of sig figs.

Chemical Reactions (Equations & Types)

  • Balanced equation: equal atoms each side; coefficients represent mole ratios.
  • Balancing steps: count atoms → adjust coefficients → re-check.
  • Five basic reaction types:
    1. Combination / Synthesis A+B\to AB.
    2. Decomposition AB\to A+B.
    3. Single-replacement A+BC\to AC+B (or non-metal Y+XZ → XY+Z).
    4. Double-replacement AB+CD\to AD+CB (precipitate, gas, or water often forms).
    5. Combustion (hydrocarbon + O₂ → CO₂ + H₂O).

Periodic Table & Periodicity

  • Arrangement by increasing Z; periodic law: properties recur periodically.
  • Atomic radius: ↓ across period; ↑ down group.
    • Cations < atoms; anions > atoms.
  • Ionization Energy IE: ↑ across period; ↓ down group.
  • Electron Affinity: more negative across; less negative down (exceptions N, O, F).
  • Electronegativity: ↑ across; ↑ up; highest F.
  • Metallic character: ↓ across; ↑ down.
  • Groups:
    • Group 1 Alkali metals +1.
    • Group 2 Alkaline earth +2.
    • Groups 3–12 Transition metals (variable charges).
    • 17 Halogens -1; 18 Noble gases (inert, full octet).
  • Special blocks: Lanthanides (4f), Actinides (5f), Metalloids along staircase (Si, Ge…).

Chemical Bonding & Electron Configurations

  • Three idealized bond types:
    1. Ionic: electrostatic attraction metal cation + non-metal anion; formula unit lattice (e.g., NaCl).
    2. Covalent: shared electron pairs between non-metals; single, double, triple bonds; follows octet rule.
    3. Metallic: lattice of cations in "sea" of delocalized e⁻ (conductivity, malleability).
  • Octet Rule: atoms gain/lose/share e⁻ to reach 8 valence electrons (He exception).
  • Electron configuration notation (Aufbau order, Pauli exclusion, Hund’s rule):
    • Example Na: 1s^22s^22p^63s^1 or abbreviated [Ne]3s^1.
  • Orbital filling order illustrated by diagonal rule diagram; remember exceptions (Cr, Cu).

Composition Stoichiometry (The Mole Concept)

  • Avogadro’s number N_A = 6.022\times10^{23}\;\text{entities mol}^{-1}.
  • Molar mass M (g·mol⁻¹) numerically equals formula weight (amu).
    • Example: H₂O → FW = 18.0\,amu → M = 18.0\,g·mol^{-1}.
  • Conversions:
    • \text{mol} = \dfrac{\text{mass (g)}}{M}\,;\; \text{entities} = \text{mol}\times N_A.
  • Percentage composition from formula:
    \%\,\text{element} = \dfrac{(n_{atoms})(AW)}{FW}\times100.
  • Empirical formula determination: convert mass % → moles → divide by smallest → use integers.
  • Molecular formula: \text{MF} = n \times \text{EF}\,;\; n = \dfrac{MW}{FW_{EF}}.
  • Example: Vitamin C 40.92%C, 4.58%H, 54.50%O → EF C3H4O3; MW 176 → n=2 → MF C6H8O6.

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