KK

Introductory Chemistry – Lecture 1: Matter, Energy, and the Scientific Method

Chemistry as the Central Science

  • Chemistry intersects with virtually every scientific and engineering discipline, providing a molecular‐level explanation for:
    • Agriculture & environmental science (e.g.
    – Nutrient cycles, soil chemistry, pollution remediation)
    • Biology & medicine (biochemical pathways, drug design)
    • Geology (mineral composition, geochemical dating)
    • Physics (thermodynamics, quantum foundations)
    • Engineering & materials science (polymers, alloys, semiconductors)
  • Practical implication: A solid foundation in chemical principles equips practitioners in diverse fields to solve real‐world problems (e.g. sustainable agriculture, renewable energy technologies).

Describing Matter: Composition & Structure

  • Two complementary descriptors:
    Composition – the kinds and proportions of atoms present.
    Structure – the spatial arrangement of those atoms.
  • "What it’s made of" + "How it’s put together" → determines properties and reactivity.

Levels of Chemical Organization

  • Atom
    • Smallest unit that retains the chemical identity of an element.
  • Element
    • Pure substance consisting of only one kind of atom.
    • Cannot be decomposed into simpler substances by ordinary chemical means.
  • Compound
    • Pure substance composed of two or more different elements chemically bonded in fixed ratios (law of definite proportions).
    • Example reaction: \mathrm{Zn}+\mathrm{S}\;\longrightarrow\;\mathrm{ZnS} (formation of zinc sulfide through chemical change).
  • Molecule
    • Discrete group of atoms held together by covalent bonds; smallest unit of a molecular compound capable of independent existence (e.g. \mathrm{H2O}, \mathrm{O2}).

Pure Substances vs. Mixtures

  • Pure substances (elements OR compounds)
    • Uniform, constant composition; single set of intrinsic properties.
  • Mixtures – physical combination of two or more substances.
    Homogeneous (solutions) – components are evenly distributed at the molecular level (e.g. salt water).
    Heterogeneous – components are unevenly distributed; phases or layers visible (e.g. sand in water, salad dressing).
  • Classification decision chart (implicit on slide):
    1. Only one type of atom? → Element.
    2. Multiple atom types but chemically bonded? → Compound.
    3. More than one substance? → Mixture.
      – Evenly blended? → Homogeneous.
      – Not evenly blended? → Heterogeneous.

Methods of Separating Mixtures (Physical, not Chemical)

  • Filtration – separates solids from liquids (sand–water).
  • Distillation – separates based on boiling‐point differences (salt water → fresh water).
  • Chromatography – separates based on differential affinities between mobile & stationary phases.
  • Magnetism, centrifugation, decanting, etc.
  • Key principle: No chemical bonds are broken; identity of each component stays intact.

States of Matter & Particle View

  • Solid – fixed shape & volume; particles closely packed, vibrate in place.
  • Liquid – fixed volume but adopts container shape; particles closely packed but can flow.
  • Gas – neither fixed shape nor volume; particles far apart, move freely.
  • Temperature is a control knob: raising T generally promotes \text{solid} \to \text{liquid} \to \text{gas} transitions.
  • Underlying axiom: "The behavior of any substance is determined by the arrangement of the particles that compose it."

Physical Properties & Physical Changes

  • Physical properties – can be measured without changing composition (mass, color, density, melting point, conductivity, magnetism).
  • Physical changes – alter state or appearance, not composition (phase changes, dissolution, cutting, grinding).
    • All phase transitions (melting, freezing, vaporization, condensation, sublimation, deposition) are physical changes.

Chemical Properties & Chemical Changes

  • Chemical properties – describe ability to undergo chemical change (flammability, acidity, oxidizing strength).
  • Chemical changes (reactions) – transform substances into new substances with different compositions & properties.
    • Example: Combination of Zn and S powders upon heating → formation of \mathrm{ZnS} (exothermic glow).
    • Indicators: color change, gas/bubble formation, precipitate, energy change, new odor.

Energy Fundamentals in Chemistry

  • Energy – capacity to do work or transfer heat.
    • Measured in joules (J) or calories (1 cal = 4.184 J).
  • Two main forms:
    1. Potential Energy (PE) – energy due to position or composition.
      – Stored in chemical bonds, gravitational height, electrostatic separation.
    2. Kinetic Energy (KE) – energy of motion.
      – E_k = \frac{1}{2} m v^2.
  • Heat (q) – energy transfer due to temperature difference; flows from higher T to lower T.

Energy & Chemical/Physical Processes

  • Systems naturally move from higher PE to lower PE when permitted – driving force for spontaneity.
  • Exothermic change – releases energy (system loses heat, surroundings warm); q<0.
  • Endothermic change – absorbs energy (system gains heat, surroundings cool); q>0.
  • Photosynthesis vs. combustion:
    • Photosynthesis stores solar energy in bonds of glucose (endothermic overall):
    6\,\mathrm{CO2}+6\,\mathrm{H2O}+\text{light} \; \longrightarrow \; \mathrm{C6H{12}O6}+6\,\mathrm{O2}.
    • Combustion of plant material releases that stored energy (exothermic).
  • "High energy" ≠ "stable"; lower‐energy arrangements are thermodynamically favored (e.g. products of exothermic reactions).

Stability, Potential Energy & Bonding

  • Bond formation lowers potential energy (release of energy → stabilization).
  • Bond breaking requires energy input (raises system energy).
  • Visual metaphor (slide): stacked blocks vs. assembled structure; stable arrangement corresponds to lower PE.

The Scientific Method: Cycle of Inquiry

  1. Make observations – qualitative & quantitative data.
  2. Formulate hypotheses – tentative explanations (testable, falsifiable).
  3. Test with experiments – controlled procedures to gather evidence.
  4. Analyze / refine – support, modify, or reject hypotheses.
  5. Iterate → accumulation of knowledge.
  • Key Terms
    Hypothesis – educated guess; narrow in scope.
    Theory – well‐substantiated, broad explanation; withstands repeated testing (e.g. atomic theory, kinetic‐molecular theory).
    Scientific law – concise statement that summarizes repeated observations (often mathematical; e.g. law of conservation of mass, PV = \text{constant} for Boyle’s law).
  • Philosophical note: Theories are not "upgraded" to laws; the two serve different roles (explanatory vs. descriptive).

Ethical, Practical & Interdisciplinary Considerations

  • Accurate classification of substances crucial for environmental monitoring (e.g. distinguishing pollutants in water matrices).
  • Understanding exothermic vs. endothermic events underpins safe industrial process design (preventing runaway reactions).
  • Scientific method fosters evidence‐based policy (e.g. climate change mitigation relies on predictive chemical models of atmospheric CO$_2$).

Summary Checklist (Self‐Test)

  • Can you differentiate element vs. compound vs. mixture?
  • Can you name two techniques for separating mixtures and explain the underlying physical property exploited?
  • Do you know all six fundamental phase changes and classify them as endo‐ or exothermic?
  • Can you articulate the distinction between a theory and a law with an example of each?
  • Can you determine whether a given process is exothermic or endothermic and predict entropy/enthalpy trends?