Is Matter Around Us Pure – Comprehensive Bullet-Point Notes

Pure Substances vs. Mixtures

• Matter can be broadly divided into:
  • Pure substances
  • Elements
    • Definition: A pure substance that cannot be split into two or more simpler substances by any chemical process.
    • Examples: Copper, oxygen, hydrogen.
    • Sub-classes:
    • Metals
      • Malleable, ductile, good conductors of heat & electricity.
      • Examples: Copper, gold, iron.
    • Non-metals
      • Non-malleable, non-ductile, poor conductors of heat & electricity.
      • Examples: Sulphur, oxygen.
    • Metalloids
      • Exhibit some metallic and some non-metallic properties.
      • Examples: Silicon, boron, arsenic.
  • Compounds
    • Definition: Pure substances made of two or more elements chemically combined in a definite ratio.
    • Key characteristics:
    • Properties differ from those of the constituent elements.
    • Possess fixed melting & boiling points.
    • Constituents can only be separated by chemical processes.
    • Examples: Water (\text{H}_2\text{O}), sugar, common salt (\text{NaCl}).
  • Impure substances (Mixtures)
  • Definition: Physical combinations of two or more elements and/or compounds in any proportion.
  • General properties:
    • Constituents retain their original properties; no new substance formed.
    • Can be separated by physical methods.
    • Do not have definite melting or boiling points.
  • Types:
    • Homogeneous mixtures (uniform composition)
    • Examples: Sugar $+$ water, water $+$ alcohol, air.
    • Heterogeneous mixtures (non-uniform composition)
    • Examples: Sand $+$ sugar, water $+$ oil.

Solutions (True Solutions)

• Definition: Homogeneous mixture of two or more substances in a single phase.
• Particle size: < 1\,\text{nm}; invisible to naked eye; do not scatter light (no Tyndall effect).
• Stability: Solute particles never settle on standing; cannot be removed by filtration.
• Components:
  • Solvent: component present in larger amount.
  • Solute: component present in smaller amount.
• Alloys
  • Mixtures of two or more metals or a metal with a non-metal (e.g.
    brass ≈ 70\% Cu + 30\% Zn).
  • Considered mixtures because they can have variable composition and exhibit properties of constituents, yet cannot be separated by ordinary physical methods.

Concentration Expressions

• Mass by mass percentage: \%\,w/w = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100
• Mass by volume percentage: \%\,w/v = \frac{\text{Mass of solute}}{\text{Volume of solution}} \times 100
• Volume by volume percentage: \%\,v/v = \frac{\text{Volume of solute}}{\text{Volume of solution}} \times 100

Solubility

• Amount of solute present in a saturated solution at a given temperature.
• Factors:
  • Temperature
  • For solids in liquids: solubility generally increases with temperature.
  • For gases in liquids: solubility decreases with rise in temperature.
  • Pressure
  • Negligible effect on solids & liquids.
  • For gases: solubility is directly proportional to pressure (Henry’s law).

Types of Solutions

• Aqueous: solvent is water (e.g.
  salt $+$ water).
• Non-aqueous: solvent is not water (e.g.
  sulphur $+$ carbon disulphide).
• Dilute vs. Concentrated: relative amount of solute dissolved.
• Unsaturated: more solute can dissolve at the given T.
• Saturated: no additional solute dissolves at the given T.

Dispersed Systems Beyond True Solutions

Suspensions

• Heterogeneous mixtures with visible particles > 100\,\text{nm} (often written >100\,\mu\text{m} in textbooks).
• Properties:
  • Particles settle on standing; mixture unstable.
  • Can be separated by filtration.
  • Scatter light and are opaque.
  • Examples: Chalk in water, wheat flour in water.

Colloids (Colloidal Solutions)

• Heterogeneous mixtures with particle size between 1\,\text{nm} and 100\,\text{nm}.
• Components:
  • Dispersed phase: colloidal particles (solute-like).
  • Dispersion medium: continuous phase (solvent-like).
• Properties:
  • Particles too small to be seen individually, yet large enough to scatter light → Tyndall effect (visible luminous path).
  • Do not settle under gravity; relatively stable.
  • Cannot be filtered by ordinary filter paper; require centrifugation or ultrafiltration.
  • Examples: Milk, fog, smoke, gelatin.

Tyndall Effect

• Phenomenon of light scattering by colloidal or very fine suspension particles, rendering the path of the beam visible.
• Absent in true solutions; present in colloids and some suspensions.
• Utilised in detecting colloidal impurities (e.g.
  in water testing) and explaining natural events (e.g.
  blue colour of sky).

Methods of Separating Mixtures

Simple Physical Methods (for heterogeneous mixtures)

• Hand-picking, Sieving, Filtration: based on particle size or visual identification.

Centrifugation

• Rapid spinning forces denser particles to bottom (sediment) and leaves lighter particles at top (supernatant).
• Used in diagnostic labs (blood/urine), dairy (cream separation).

Separating Funnel (Immiscible Liquids)

• Separates liquids forming distinct layers due to density difference.
• Applications: Oil–water separation; extraction of iron from ore (removal of slag).

Evaporation

• Separates a volatile solvent from a non-volatile solute.
• Example: Recovering coloured dye from ink.
• Limitations: May decompose or char heat-sensitive solids.

Sublimation

• Separates sublimable solids (ammonium chloride, camphor, naphthalene, anthracene) from non-sublimable impurities.

Chromatography

• Based on differential adsorption and solubility of components on/in a stationary & mobile phase.
• Separates mixtures of coloured substances (e.g.
  plant pigments, pen inks).
• More soluble component travels faster and hence separates first.

Distillation (Simple)

• Used for two miscible liquids with large boiling point difference (≥ 25\,\text{K}) that do not decompose.
• Example: Acetone–water.

Fractional Distillation

• Employed when the B.P. difference is < 25\,\text{K} or for multiple liquids.
• Set-up uses a fractionating column to provide repeated condensation–evaporation cycles, enhancing separation.
• Industrial example: Separating liquid air
  1. Compress \& cool air → liquefy.
  2. Fractionally distil; gases vaporise at specific column heights (e.g.
     \text{N}2 at 77\,\text{K}, \text{O}2 at 90\,\text{K}, Ar in between).

Crystallization

• Purifies solids by forming crystals after partial evaporation or cooling of a hot saturated solution.
• Advantages over simple evaporation:
  1. Heat-labile solids (e.g.
     sugar) may char on strong heating.
  2. Dissolved impurities that escape filtration remain after evaporation but exclude themselves from crystals.
• Industrial use: Sugar refining, production of alum.

Physical vs. Chemical Changes

• Physical Change
  • Alters physical properties (state, colour, hardness, density, M.P., B.P.) without forming new substances.
  • Usually reversible; energy change minimal.
• Chemical Change
  • Produces one or more new substances with different chemical properties (odour, combustibility, reactivity).
  • Generally irreversible; involves significant energy change (heat, light, sound).

Inter-connections & Significance

• Understanding purity is essential in pharmaceuticals, food safety, metallurgy, and environmental monitoring.
• Techniques like fractional distillation of air supply industrial \text{O}2/\text{N}2 critical for medicine & metallurgy.
• Colloid science underpins products from paints to medicines (e.g.
  drug delivery suspensions).
• Chromatography is foundational for forensic science, drug testing, and biochemical research, stressing ethical handling of data.
• Awareness of physical vs. chemical changes helps prevent hazards (e.g.
  flammability in chemical storage) and guides sustainable material usage.