Chapter 2 — Is Matter Around Us Pure

Overview

Matter can be either a pure substance or a mixture. A pure substance has a fixed, definite composition and distinct properties (fixed melting and boiling points). A mixture contains two or more substances physically combined in variable proportions and the components retain their own properties.

Pure substances

• Element: Made of only one type of atom (e.g., oxygen, iron).
• Compound: Made of two or more elements chemically combined in a fixed ratio (e.g., water H2O, sodium chloride NaCl).
• Key property: Sharp/fixed melting and boiling points; cannot be separated by physical methods.

Mixtures

• Definition: Physical combination of two or more substances with variable composition.
• Classification by uniformity:
  • Homogeneous mixture (solution): Composition is uniform throughout (e.g., sugar solution, air). True solutions have particle size < 1 nm and do not show Tyndall effect.
  • Heterogeneous mixture: Non-uniform composition (e.g., soil, salad). Components often visible or separable by simple methods.
• Classification by particle size:
  • Solution: < 1 nm (no scattering of light).
  • Colloid: about 1–1000 nm (shows Tyndall effect; shows Brownian motion; e.g., milk, fog).
  • Suspension: > 1000 nm (particles settle on standing; e.g., muddy water).

Distinguishing features

• Pure substances: fixed composition; sharp melting/boiling point; properties are constant.
• Mixtures: variable composition; melting and boiling over a range (impure substances melt/boil over ranges); components retain properties and can be separated by physical means.

Common separation techniques (when to use)

• Filtration: Separate insoluble solids from liquids (suspension).
• Sedimentation and decantation: Let heavy particles settle then pour off liquid (suspension).
• Evaporation and crystallization: Recover dissolved solids from solution (e.g., obtain salt from salt solution).
• Distillation: Separate miscible liquids with different boiling points or to purify a liquid (simple and fractional distillation).
• Fractional distillation: Separate miscible liquids with closer boiling points (e.g., fractional distillation of petroleum fractions, separation of components of air by low-temperature fractional distillation).
• Chromatography: Separate components of a small mixture based on different affinities (useful for inks, pigments).
• Sublimation: Separate substances that sublime (solid to gas) from non-subliming solids (e.g., dry ice from mixture under appropriate conditions).
• Centrifugation: Rapidly separate fine particles from liquids by spinning (used in laboratories, blood separation).
• Magnetic separation: Separate magnetic substances from non-magnetic (e.g., iron filings from sand).

Special effects and observations

• Tyndall effect: Scattering of light by colloidal particles — visible light beam in a colloid but not in a true solution.
• Brownian motion: Random motion of colloidal particles caused by collisions with solvent molecules — evidence for molecular motion.

Why this matters

Recognizing pure substances and mixtures and knowing correct separation techniques are essential in chemistry, industry (purification, metallurgy), environmental science (water treatment), and everyday life (making beverages, cleaning, medicine preparation).

Definition: A mixture is a combination of two or more substances that are not chemically combined and can be separated by physical methods. Mixtures are classified on the basis of uniformity (homogeneous or heterogeneous) and on the basis of particle size (solutions, colloids, suspensions).

1. By uniformity

• Homogeneous mixtures: Uniform composition throughout; single visible phase. Examples: sugar dissolved in water, alloys, air. Properties: cannot be separated by simple filtration; no visible particles.
• Heterogeneous mixtures: Non-uniform composition; two or more visible phases. Examples: soil, salad, oil + water. Properties: components can often be seen and separated by physical methods like filtration, decantation.

2. By particle size — this is a very useful classification for physical behaviour and separation method:

• Solutions (true solutions): Particle size < 1 nm. Particles (ions or molecules) are too small to scatter light (no Tyndall effect), do not settle on standing, and cannot be separated by ordinary filtration. Example: salt or sugar in water, air.
• Colloids (colloidal solutions): Particle size ≈ 1 nm – 1000 nm (1 nm – 1000 nm is commonly used). Particles are intermediate in size; they scatter light (Tyndall effect), do not settle on standing, and cannot be separated by ordinary filter paper but can be separated by ultrafiltration or centrifugation. Examples: milk (liquid emulsion), fog (aerosol), sols like paint.
• Suspensions: Particle size > 1000 nm (> 1 μm). Particles are large enough to scatter light strongly, they settle on standing, and can be separated by filtration. Example: muddy water, sand in water.

Key distinguishing properties (solutions / colloids / suspensions):

• Tyndall effect: solutions — no; colloids — yes; suspensions — yes (but due to large particles).
• Stability: solutions — very stable; colloids — kinetically stable (do not settle quickly); suspensions — unstable (settle on standing).
• Filterability: solutions — pass through filter paper; colloids — pass through ordinary filter paper but not through semipermeable membranes/ultrafilters; suspensions — retained by filter paper.

Common separation techniques and when to use them:

• Filtration: separate insoluble solids from liquids (suspensions).
• Sedimentation and decantation: for suspensions where particles settle by gravity.
• Centrifugation: speeds up settling for fine particles (colloids approaching suspension behaviour).
• Evaporation / crystallization: to recover dissolved solids from solutions.
• Distillation (simple and fractional): separate components of a homogeneous liquid mixture based on boiling points (solutions of volatile liquids).
• Chromatography: separate components of a homogeneous mixture based on affinity for stationary/mobile phases (useful for coloured mixtures and analysis).
• Sublimation: separate a sublimable solid from non-sublimable impurities.
• Magnetic separation, sieving: for special heterogeneous mixtures (e.g., iron filings, different particle sizes).

Why this classification matters: The classification predicts physical behaviour (visibility, settling, light scattering) and tells which separation method to use. It is foundational for laboratory techniques, environmental science (air/water quality), food science (emulsions), and industrial processes (alloys, suspensions).

What is a solution? A solution is a homogeneous mixture of two or more substances. It has a uniform composition throughout and consists of a solute (present in smaller amount) dissolved in a solvent (present in larger amount).

• Solute: The substance that gets dissolved (e.g., salt, sugar, CO2).
• Solvent: The medium that dissolves the solute (commonly water for aqueous solutions; can be ethanol, benzene, etc.).
• Types of solutions (by physical state): solid in liquid (salt in water), gas in liquid (CO2 in soda), liquid in liquid (alcohol in water), solid in solid (alloys like brass).

Concentration — meaning and qualitative terms
Concentration tells how much solute is present in a given amount of solution. Common qualitative terms: dilute (little solute) and concentrated (large amount of solute). Solutions can also be saturated (no more solute dissolves at given T), unsaturated (more solute can dissolve) or supersaturated (contains more solute than normally possible at that T).

How concentration is changed

• Add or remove solute (e.g., add sugar to increase concentration).
• Add or remove solvent (e.g., add water to dilute).
• Change temperature for many solids in liquids (heating usually increases solubility, cooling decreases it).

Example calculation (mass percent)
If 5 g of salt is dissolved in 95 g of water, mass of solution = 5 + 95 = 100 g. Mass percent of salt = (5/100) × 100 = 5% w/w.

Importance / real life: Concentration is important in medicine (drug doses), cooking (taste), industry (electroplating, chemical manufacture), environmental science (pollutant levels like ppm), and everyday uses like making tea/coffee or a saline solution.

Definition and particle-size classification

• Solution (true solution): A homogeneous mixture in which solute particles are individual molecules or ions. Particle size < 1 nm (10⁻⁹ m). Examples: sugar in water, salt solution.
• Colloid: A heterogeneous mixture where dispersed particles (called dispersed phase) have sizes large compared to molecules but small enough to remain suspended. Typical particle size is about 1 nm to 1000 nm (10⁻⁹ to 10⁻⁶ m). Colloids do not settle on standing and cannot be separated by ordinary filter paper.
• Suspension: A heterogeneous mixture with relatively large particles (>1000 nm or >1 μm). Particles scatter light, settle on standing, and can be separated by filtration.

Key properties that distinguish them

• Tyndall effect: Colloidal particles scatter light, making a light beam visible when passed through the mixture (Tyndall effect). True solutions do not show this; suspensions may scatter strongly but particles are large enough to settle.
• Filtration and settling: Suspensions settle on standing (gravity causes sedimentation) and can be removed by ordinary filtration. Colloids are stable and do not settle; they pass through ordinary filter paper but can be separated only by ultrafiltration or coagulation.
• Brownian motion: Colloidal particles undergo random zig-zag motion due to collisions with solvent molecules; this motion helps keep colloids dispersed. Brownian motion is not observable in true solutions or in large suspended particles.
• Stability: Colloids are relatively stable (may coagulate under certain conditions). Suspensions are unstable and separate with time.

Why these differences arise

Differences come from particle size and interaction with the medium. Very small particles (solutions) are molecularly dispersed. Colloidal particles are large enough to scatter light and experience continual bombardment by solvent molecules (Brownian motion), preventing settling. Very large particles in suspensions are affected more by gravity and overcome random thermal motion, so they settle.

Practical notes for identification

• Shine a flashlight: visible beam in a colloid (Tyndall effect), not in a true solution.
• Let the mixture stand: suspensions show settling, colloids do not (unless coagulated), true solutions show no settling.
• Filter with ordinary filter paper: suspensions leave residue, colloids pass through.

Introduction
A mixture contains two or more substances physically combined. Mixtures are either homogeneous (uniform composition, e.g., salt solution) or heterogeneous (non-uniform, e.g., sand + water). Separation techniques use physical properties (particle size, solubility, boiling point, magnetic property, volatility, etc.) to separate components without chemical change.

Common separation techniques (principle, method, use)

• Filtration — separates an insoluble solid from a liquid (heterogeneous). Principle: particle size. Method: pour mixture through filter paper/cloth; liquid (filtrate) passes, solid (residue) stays. Example: sand from water.
• Sedimentation and Decantation — allow heavy particles to settle (sedimentation), then pour off the clear liquid (decantation). Used for mixtures with large density/size differences. Example: muddy water.
• Evaporation and Crystallization — remove solvent by heating to get dissolved solute (evaporation) or heat a saturated solution then cool to form pure crystals (crystallization). Principle: difference in volatility and solubility with temperature. Example: obtaining salt from seawater (evaporation); purifying salt by crystallization.
• Simple Distillation — separates a liquid from non-volatile solute or liquids with very different boiling points. Heat mixture, vaporise the lower-boiling component, condense vapor to collect distillate. Example: separating water from salt solution.
• Fractional Distillation — separates two miscible liquids with relatively close boiling points using a fractionating column to achieve repeated vaporisation–condensation cycles. Example: separation of alcohol and water in laboratory (industrial: crude oil fractionation).
• Chromatography — separates components based on differing affinities between stationary phase and mobile phase. Types include paper chromatography, column chromatography. Method: sample placed on stationary phase; solvent (mobile phase) moves and carries components at different rates. Example: separating plant pigments; checking purity of substances.
  Key concept: Rf value (relative mobility) = distance moved by substance / distance moved by solvent front.
• Sublimation — separates a mixture where one component sublimes (solid <> gas) and the other does not. Heat; sublime the volatile solid, condense the vapour on a cool surface. Example: separating iodine from salt or ammonium chloride from sand.
• Magnetic separation — uses magnetic properties to remove magnetic materials from mixtures. Example: iron filings removed from sand using a magnet.
• Sieving (screening) — separates particles by size using mesh/sieve. Example: separating stones from grain.
• Winnowing — uses wind or air blow to separate lighter (chaff) from heavier (grain) particles. Example: separating rice/wheat from chaff.
• Centrifugation — uses rapid spinning to separate components of different densities (e.g., suspended solids from liquid). Principle: produces large effective gravity (centrifugal force) to speed sedimentation. Example: separating cream from milk in dairies; lab separation of blood components.
• Solvent Extraction — uses difference in solubility in two immiscible liquids to transfer a substance into the solvent in which it is more soluble. Example: extracting oil from seeds using an organic solvent.

Choosing a method
Select a technique based on the physical properties that differ among components: particle size (filtration, sieving), volatility (distillation, evaporation, sublimation), solubility (crystallization, solvent extraction), magnetic property (magnetic separation), affinity for stationary/mobile phases (chromatography), or density (centrifugation, decantation).

Advantages & Limitations (brief)

• Most techniques are simple and reversible (no chemical change).
• Purity depends on technique and conditions (e.g., fractional distillation gives better separation for close boiling points than simple distillation).
• Some methods need special apparatus or are energy intensive (distillation, crystallization).

Safety & Practical Tips: Wear protective eyewear when heating or using chemicals; ensure good ventilation for volatile solvents; use appropriate glassware for distillation.

What is purity? A pure substance contains only one type of particle (element or compound) and has fixed characteristic properties (melting point, boiling point, density, etc.). A mixture contains two or more substances physically combined and shows variable properties.

Why test for purity? Identifying whether a sample is pure is important in laboratories, industry, medicine and food safety because impurities change physical and chemical behaviour and may be harmful or reduce effectiveness.

Main tests of purity (Class 9 level)

• Melting point and boiling point: A pure solid melts at a specific, sharp temperature; a pure liquid boils at a sharp temperature under fixed pressure. Impurities usually lower the melting point and broaden the melting range, and they generally raise the boiling point of a liquid (relative to the pure substance). Measuring these gives a quick test of purity.
• Paper chromatography: A technique to separate and identify components of a mixture (especially coloured substances). A spot of the mixture is placed on chromatography paper; the paper is dipped in a solvent which rises by capillary action and carries components at different rates. Pure substances give a single spot (or single Rf value); mixtures give two or more spots.
• Solubility, filtration and crystallisation: Differences in solubility are used to separate and test purity. Complete dissolution followed by recrystallisation can purify a solid. Impure samples often form crystals of different appearance and may contain insoluble residues on filtration.
• Conductivity and colour/turbidity: For some samples (e.g., water) electrical conductivity, turbidity or colour changes indicate dissolved ionic or particulate impurities.

How impurities affect measurements: Impurities disrupt the regular lattice of a solid (lowering melting point and giving a broad melting range). In liquids, solutes lower vapour pressure, which causes boiling point elevation (requires more heat to boil) — a concept used qualitatively in purity testing.

Practical procedure (paper chromatography - brief):

1. Draw a pencil baseline near the bottom of chromatography paper and spot the sample on it.
2. Place paper in a container with a small amount of solvent so the solvent level is below the spot.
3. Allow solvent to rise until near the top; remove and mark solvent front immediately. Let paper dry and observe separated spots.
4. Calculate Rf values for identification (distance of spot from baseline divided by distance of solvent front from baseline).

Interpretation: If a sample shows a single sharp melting point equal to the expected value, or a single chromatographic spot (with Rf matching a known substance), it is likely pure. Multiple spots, depressed/ broadened melting range, unexpected conductivity or colour indicate impurities.

Limitations: Very small amounts of impurity may not change melting/boiling points much. Chromatography separates components but identification requires comparison with standards.