Chemistry – Matter Classification, Properties, States & Temperature

Classification of Matter

• Matter

  • Defined as anything that has mass and occupies space.
  • Makes up all commonly-used materials (e.g., water, wood, plastic bags).

• Overview of Classifications

  • Two fundamental categories based on composition:
  • Pure substances – fixed/definite composition.
  • Mixtures – physical blends of ≥2 different substances, variable composition, not chemically bonded.

Pure Substances

• General Characteristics
  • One unique, constant composition throughout the sample.
  • Always homogeneous at the atomic/molecular level.
  • Further divided into elements and compounds.

Elements

• Contain only one kind of atom.
• Cannot be decomposed by ordinary chemical means.
• Examples & symbols: copper (Cu), lead (Pb), aluminum (Al).
• Microscopic representation: every particle in an Al can is an Al atom (identical).
• Practical importance: provide building blocks for all other matter.

Compounds

• Composed of two or more elements chemically combined in a fixed ratio.
• Exhibit properties different from their constituent elements.
• Chemical formulas specify the ratio.
  • Hydrogen peroxide: \text{H}2\text{O}2 (2 H : 2 O).
  • Water: \text{H}_2\text{O}.
  • Sodium chloride (table salt): \text{NaCl}.
  • Sucrose (table sugar): \text{C}{12}\text{H}{22}\text{O}_{11}.
• Decomposition reaction illustration: \text{2 NaCl} \rightarrow 2\,\text{Na} + \text{Cl}_2 (produces elemental sodium & chlorine).

Mixtures

• Key Traits
  • Physical combination → each component retains its identity and properties.
  • Components present in variable proportions.
  • Can be separated by physical techniques (filtration, distillation, chromatography, decanting, straining, etc.).

Homogeneous Mixtures (Solutions)

• Uniform composition; visually appear as one phase.
• No visible boundaries between components.
• Laboratory & everyday examples:
  • Brass (Cu + Zn).
  • Scuba breathing gases:
  • Nitrox – O₂ + N₂.
  • Heliox – O₂ + He.
  • Trimix – O₂ + He + N₂.

Heterogeneous Mixtures

• Composition varies from one region to another; distinct phases visible.
• Examples: copper metal chunks in water, orange juice with pulp, spaghetti in water (requires strainer to separate).

Physical Separation Methods

• Filtration – separates solid–liquid mixtures (coffee filter, lab funnel).
• Chromatography – components travel at different rates on a stationary phase (ink pigments on paper, drug analysis).
• Straining/Decanting – coarse mechanical separation (pasta-water).

Classification Practice (Sample Exercise #1)

• Wine → Homogeneous mixture.
• Gold → Element.
• \text{CO}_2 → Compound.
• Orange juice with pulp → Heterogeneous mixture.

Physical States of Matter

• Matter exists primarily as solids, liquids, gases.

Solids

• Definite shape & volume.
• Particles tightly packed in rigid lattice, vibrate slowly.
• Very strong intermolecular/ionic/metallic interactions.
• Examples: ice, salt, iron, amethyst (purple \text{SiO}_2 quartz).

Liquids

• Definite volume but no definite shape (take container shape).
• Particles close but can slide/flow; moderate interactions.
• Examples: water, oil, vinegar, eye drops.

Gases

• Neither definite shape nor volume; expand to fill container.
• Particles far apart, very fast, negligible interactions.
• Examples: water vapor, helium, air.

Comparative Summary (Table 3.1)

• Shape: solid – fixed, liquid – container, gas – container.
• Volume: solid & liquid – fixed, gas – fills container.
• Particle arrangement: solid – fixed/close; liquid – random/close; gas – random/far.
• Interaction strength: very strong → virtually none.
• Motion: very slow → very fast.

Identifying States (Learning Check)

• Vitamin tablets → solid.
• Eye drops, vegetable oil → liquids.
• Candle (wax) → solid (until melted).
• Air in basketball → gas.

Physical vs. Chemical Properties & Changes

Physical Properties

• Observed/measured without altering chemical identity.
• Examples: shape, phase, color, density, melting/freezing/boiling points.
• Copper specifics: reddish-orange, shiny, excellent conductor, \text{m.p.}=1083\,^\circ\text{C}, \text{b.p.}=2567\,^\circ\text{C}.

Physical Changes

• Change in form or state, composition constant.
• Phase changes: ice ↔ water ↔ steam.
• Dissolving, cutting, hammering gold into foil.

Chemical Properties

• Describe a substance’s ability to form new substances.
• Require chemical change to observe (flammable, rusting, reactivity with acid).

Chemical Changes

• Produce substances with new compositions & properties.
• Evidence: color change, gas formation, precipitate, heat/light.
• Examples:
  • Iron + \text{O}2 → rust \text{Fe}2\text{O}_3 (red-orange).
  • Sugar caramelizing upon heating (complex polymerization & dehydration).

Physical vs Chemical Examples (Table 3.3)

• Physical: boiling water, drawing copper wire, dissolving sugar, cutting paper.
• Chemical: silver tarnishing, wood burning, heating sugar to caramel, iron rusting.

Summary Matrix (Table 3.4)

• Property: physical vs chemical (ability to change).
• Change: physical vs chemical (actual transformation).

Characterizing a Mystery Metal (Sample Exercise #2)

• Physical: silvery white, lustrous; melts at 649\,^\circ\text{C}; boils at 1105\,^\circ\text{C}; density 1.738\,\text{g/cm}^3; malleable/ductile; good electrical conductor.
• Chemical: burns in air with intense white light; reacts with \text{Cl}_2 to form brittle white solid.

Change Identification Practice (Sample Exercise #3)

• Rusting iron → chemical.
• Sugar dissolving → physical.
• Burning log → chemical.
• Melting ice → physical.
• Grinding cinnamon → physical (size reduction).

Temperature

Concept

• Indicates relative hotness/coldness vs reference.
• Measured with a thermometer.
• Scientific standard: degrees Celsius (°C).

Scales Compared

• Celsius (°C), Fahrenheit (°F), Kelvin (K).
• Reference: water freezes/boils.
  • Celsius: 0\,^\circ\text{C} → freeze; 100\,^\circ\text{C} → boil.
  • Fahrenheit: 32\,^\circ\text{F} → freeze; 212\,^\circ\text{F} → boil.
  • Kelvin: 273\,\text{K} → freeze; 373\,\text{K} → boil.
• Interval equivalence: 180\,\text{°F}=100\,\text{°C}.

Conversions

• T{\text{F}} = 1.8\,T{\text{C}} + 32.
• T{\text{C}} = \dfrac{T{\text{F}} - 32}{1.8}.
• T{\text{K}} = T{\text{C}} + 273 (exact conversion; 1 K = 1 °C).

Worked Example: Room Temperature

• Given 21\,^\circ\text{C}; find T{\text{F}}. T{\text{F}} = 1.8\times21 + 32 = 70\,^\circ\text{F}.

Kelvin Scale & Absolute Zero

• Coldest possible temperature = -273\,^\circ\text{C}=0\,\text{K} (no negative K).
• Same unit size as °C but no degree symbol.

Practice Problems

• Normal body temp 37\,^\circ\text{C} → T_K = 37+273 = 310\,\text{K}.
• Cold winter day -15\,^\circ\text{F} → T_C = \dfrac{-15 - 32}{1.8} \approx -26\,^\circ\text{C} (rounded to ones place).
• Hypothermia case 34.8\,^\circ\text{C} → T_F = 1.8(34.8)+32 = 94.6\,^\circ\text{F}.

Comparative Table (3.5)

• Highlights key environment/health benchmarks, e.g., sun’s surface ≈ 9937\,^\circ\text{F}, nitrogen liquefies ≈ -346\,^\circ\text{F}, absolute zero -459\,^\circ\text{F}.

Health Connections

• Hyperthermia (body T > 41\,^\circ\text{C}): may cause convulsions, brain damage; treat with ice-water immersion.
• Heatstroke: threshold \ge 41.1\,^\circ\text{C}.
• Hypothermia (body T ≈ 28.5\,^\circ\text{C}): treat with warmed O₂, fluids, peritoneal lavage at 37\,^\circ\text{C}.

Integration & Real-World Significance

• Accurate classification of matter aids in predicting behavior (e.g., choosing suitable materials for scuba tanks → homogeneous gas mixtures prevent variable partial pressures).
• Understanding physical vs chemical changes critical for recycling (physical separation vs chemical reprocessing).
• Temperature mastery essential in laboratory safety, medical diagnostics (fever vs hypothermia), industrial control (boiling points guide distillation).
• Ethical & practical considerations: correct handling of hazardous compounds (sodium metal vs NaCl), environmental relevance (rusting infrastructure, corrosion inhibitors).