ACCL Chemistry – Chapter 1: Matter and Scientific Method

• Week-1 mastery goals (due 19 July 2021)
  • Math & Physics \to 100\%
  • Measurement \to 40\%
  • Matter \to 60\%
  • Advanced topic pH / Acid–Base at 31\% (stretch goal)

Introduction to Chemistry

• Working definition: Chemistry is the study of matter, its properties, the changes it undergoes, and the energy associated with those changes.

• Five classical branches (historical grouping)

  1. Organic Chemistry

  2. Inorganic Chemistry

  3. Physical Chemistry

  4. Analytical Chemistry

  5. Biochemistry
     • Modern sub-disciplines (examples enumerated on slides): environmental, materials, polymer, medicinal, nuclear, theoretical, forensic, astro-, green chemistry, etc.

• Importance: chemistry bridges physics (microscopic forces/energetics) and biology (macroscopic living systems); provides foundation for engineering, pharmacology, environmental science.

Key Ideas & Vocabulary Around Matter

Fundamental Definition

• Matter = anything that possesses both mass and volume (occupies space).

Physical vs. Chemical Properties

• Property = any characteristic that can be used to identify or describe matter.

• Physical properties: observed without altering chemical identity. Examples: melting point, boiling point, density, color, hardness, conductivity, taste, texture, mass, volume, shape.

• Chemical properties: can only be observed via a chemical change (new substances formed). Examples: flammability, acidity, basicity, corrosiveness, oxidizing ability, combustibility, explosiveness, stability, reactivity, rusting tendency.

Intensive vs. Extensive

• Intensive: independent of sample size; e.g., temperature, density, color.

• Extensive: depends on amount present; e.g., mass, volume, heat.
  • Illustration: Temperature (intensive) vs. Heat (extensive).

States of Matter & Macroscopic Properties

• Fixed = property unchanged by container; Indefinite = adopts container attribute.

• Structure determines properties: arrangement and freedom of atoms/molecules differ among the three states, explaining mechanical properties, compressibility, and flow behavior.

Compressibility Illustration

• Gases: large intermolecular spacing ⇒ highly compressible (Fig 3.7).

• Solids & liquids: particles already close ⇒ essentially incompressible.

Solids – Crystalline vs. Amorphous

• Crystalline: long-range ordered lattice (e.g., NaCl, diamond).

• Amorphous: no extended periodic order (e.g., glass, most plastics).

Phase Changes (Physical Transitions)

• Melting: solid → liquid

• Freezing: liquid → solid

• Boiling/Vaporization: liquid → gas

• Condensation: gas → liquid

• Sublimation: solid → gas

• Deposition: gas → solid

• Note: Evaporation discussed as a solution process, not purely a bulk phase change.

Physical vs. Chemical Change

• Physical change: different form of same substance; molecular identity unchanged (e.g., dissolving, cutting, phase transitions).

• Chemical change: results in new substances with new molecular identities; evidences: color change, gas evolution, heat/light, precipitate formation.

Practice Identifications (selected slide answers)

• Salt white/granular → physical property.

• Salt melts at 801\,^{\circ}\text{C} → physical.

• Stability at room T → chemical property (no decomposition).

• Solubility 36\,\text{g/100 g H}_2\text{O} → physical.

• Conductivity of molten/aqueous salt → physical.

• \text{AgNO}_3 + brine → white precipitate (AgCl) → chemical property/change.

• Electrolysis of molten salt → Na metal + Cl2 gas → chemical change.

Classification of Matter

Pure Substances vs. Mixtures

• Pure substance: constant composition; identical pieces in identical % composition; cannot be separated by physical means; temperature remains constant during melting/boiling.
  • Subcategories: Elements and Compounds.

• Mixture: variable composition; components physically combined; separable via physical techniques; melting/boiling ranges vary with composition.
  • Homogeneous (solution): uniform; appears single phase.
  • Heterogeneous: visibly distinct phases/regions.

Pure Substances – Definitions

• Element: only one kind of atom; \sim 91 naturally occurring (of 116 known). Oxygen most abundant by mass in Earth’s crust & human body.

• Compound: chemical combination of two + elements; decomposable by chemical means.

• Molecule: smallest discrete unit retaining chemical properties

Mixtures – Separation Techniques (physical)

Illustrative Set-ups

• Distillation apparatus: boiling flask, condenser with water jacket, collection flask ⇒ more volatile component condenses first.

• Gravity/Vacuum Filtration: funnel + filter paper retains solid; filtrate collected.

Conservation of Mass & Early Scientific Debate

• Law of Conservation of Mass (Lavoisier, 1789):
  • “Matter is neither created nor destroyed in a chemical reaction.”
  • \sum m{\text{reactants}} = \sum m{\text{products}}.

• Historical context: disproved Phlogiston theory (imaginary fire-substance); ushered quantitative methods.

Worked Mass Balance Example

• Butane combustion:
  \text{butane} + \text{O}2 \to \text{CO}2 + \text{H}2\text{O} Data: 58\,\text{g} butane + 208\,\text{g} \text{O}2 → 176\,\text{g} \text{CO}2 + 90\,\text{g} \text{H}2\text{O}
  58 + 208 = 266 g = 176 + 90 = 266 g ✔️

Practice: Sugar + H$2$SO$4$ Demo

• Initial combined mass: 144.0\,\text{g}.

• Final beaker + carbon snake mass: 129.6\,\text{g}.

• Steam (gaseous H$_2$O) lost:
  144.0\,\text{g} - 129.6\,\text{g} = 14.4\,\text{g}.

The Scientific Method in Chemistry

1. Observation / Problem Statement
   • Collection of qualitative & quantitative data.

2. Hypothesis – tentative, testable explanation.

3. Experimentation – controlled tests; gather new data to challenge hypothesis.

4. Conclusion
   • If repeatedly validated ⇒ becomes Theory (well-substantiated explanatory model).
   • Broad, consistent observations summarized as Law (statement of ‘what’ happens).

Theory vs. Law (slide MC question)

• Correct distinction: (b) A law summarizes a series of related observations, while a theory gives the underlying reasons for them.

Connection to Prior & Future Content

• Measurement section (next lecture) will formalize units, significant figures—critical for quantitative conservation laws.

• Atoms/Ions/Molecules & Stoichiometry chapters will build molecular-level basis underpinning macroscopic mass relationships.

• Thermochemistry links energy changes (1st Law) to matter changes; complements conservation of mass.

• Ethical / practical relevance:
  • Mass balance necessary for industrial scale-up, environmental emissions auditing.
  • Scientific method underlies evidence-based policy and debunking pseudoscience (e.g., phlogiston, modern misinformation).