Ch 02 - IS MATTER AROUND US PURE?9th Science

A pure substance consists of a single type of particle (atoms or molecules). All constituent particles have the exact same chemical nature. They cannot be separated into other kinds of matter by any physical process.

Examples: Gold (Au), pure water ($H_2O$), sucrose.

An element is a basic form of matter composed entirely of one kind of atom. It cannot be broken down into simpler substances by ordinary chemical reactions (such as applying heat, light, or electrical energy).

Historical Note: Antoine Lavoisier established the first experimentally useful definition of an element.

Standard Liquids: Only two elements are liquid at standard room temperature (298 K): Mercury (Hg) and Bromine (Br).

Near-Room Temperature Liquids: Gallium (Ga) and Cesium (Cs) become liquid at a temperature slightly above room temperature (approx. 303 K or 30 degrees C).

Metals: Lustrous (shiny), malleable (can be beaten into thin sheets), ductile (can be drawn into wires), sonorous (ring when struck), and good conductors of heat and electricity.

Non-Metals: Display a variety of colours, are non-lustrous, brittle (not malleable or ductile), non-sonorous, and poor conductors.

Key Exceptions:

1. Graphite (a non-metal carbon allotrope) conducts electricity.
2. Iodine (a non-metal) is lustrous.
3. Mercury is a metal but exists as a liquid at room temperature.

Metalloids are elements that exhibit intermediate properties, falling right between those of metals and non-metals. Because of these intermediate properties, they often act as semiconductors.

Key Examples: Boron (B), Silicon (Si), and Germanium (Ge).

A mixture contains more than one pure substance (elements or compounds) mixed in any proportion.

Unlike pure substances, it has a variable composition. A critical distinction is that a mixture strictly retains the chemical properties of its individual components. These components can be separated fairly easily by physical methods (e.g., filtration, distillation).

Variables & Units:

• Mass of Solute = Mass of the dissolved substance (g or kg)
• Mass of Solution = Mass of Solute + Mass of Solvent (g or kg)

Conditions/Notes:

• Ensure both masses are in the same SI or metric unit before calculating.
• Often used for solid-in-liquid or solid-in-solid solutions.

Homogeneous Mixtures:

• Have a uniform composition throughout the mass.
• There are no visible boundaries of separation between components.
• Examples: Saltwater, air, vinegar, alloys.

Heterogeneous Mixtures:

• Have a non-uniform composition.
• Visible boundaries of separation exist between physically distinct parts.
• Examples: Sand and salt, oil and water, blood, wood.

Variables & Units:

• Mass of Solute = Mass of the dissolved substance (g)
• Volume of Solution = Total volume of the final solution (mL or L)

Conditions/Notes:

• Very commonly used in pharmacology and daily laboratory preparations (e.g., 5% w/v saline solution).

Compounds:

• Elements react chemically to form entirely new substances.
• The composition is always fixed (a definite ratio by mass).
• Properties are totally different from the constituent elements.
• Can only be separated by chemical or electrochemical reactions.

Mixtures:

• Substances mix physically without forming new substances.
• Have a variable composition.
• Retain the properties of their constituents.
• Separated easily by physical methods.

An alloy (e.g., Brass = 30% Zn + 70% Cu) is a solid homogeneous mixture of two or more metals, or a metal and a non-metal.

It is scientifically classified as a mixture because:

1. It shows the properties of its constituent elements.
2. It can have a variable composition (e.g., the exact ratio of copper to zinc can vary in different types of brass).

1. Homogeneity: It is a perfectly homogeneous mixture.
2. Particle Size: Solute particles are extremely small, < 1 nm ($10^{-9}$ m) in diameter. They cannot be seen by the naked eye.
3. Tyndall Effect: Because of the minute particle size, they do not scatter light. The path of light is invisible.
4. Stability: It is highly stable; solute particles do not settle down when left undisturbed and cannot be separated by ordinary filtration.

A suspension is a heterogeneous mixture where solid particles are dispersed in a liquid but remain undissolved.

• Particle Size: Large enough to be seen by the naked eye (> 100 nm).
• Tyndall Effect: Suspended particles scatter a beam of light passing through it, making its path visible.
• Stability: It is unstable. Particles settle down due to gravity when left undisturbed. Once settled, the suspension breaks and no longer scatters light.

A colloid (like milk) is actually a heterogeneous mixture consisting of distinct phases.

• Particle Size: Intermediate size (1 nm to 1000 nm). They are too small to be seen individually, causing the mixture to incorrectly look homogeneous.
• Tyndall Effect: Particles are large enough to strongly scatter visible light.
• Separation: Quite stable (they do not settle). They cannot be separated by ordinary filtration, but can be separated using centrifugation.

1. Dispersed Phase (The 'Solute'): The solute-like component or the active particles that are distributed throughout the mixture.
2. Dispersion Medium (The 'Solvent'): The continuous medium or substance in which the dispersed phase is suspended.

Example: In fog, liquid water droplets (dispersed phase) are suspended in air (dispersion medium).

Aerosol: Liquid or Solid dispersed in Gas (e.g., Fog, Smoke, Automobile exhaust).
Foam: Gas dispersed in Liquid or Solid (e.g., Shaving cream, Sponge, Pumice stone).
Emulsion: Liquid dispersed in Liquid (e.g., Milk, Face cream).
Sol: Solid dispersed in Liquid (e.g., Milk of magnesia, Mud).
Solid Sol: Solid dispersed in Solid (e.g., Coloured gemstones, Milky glass).
Gel: Liquid dispersed in Solid (e.g., Jelly, Cheese, Butter).

Saturated Solution: A solution that has dissolved as much solute as it is capable of dissolving at a given temperature. No more solute can dissolve without altering conditions.

Solubility: The maximum amount of solute present in a saturated solution at a specific temperature.

Temperature Effect: Heating increases the kinetic energy of the solvent. This increases solubility. A saturated solution becomes unsaturated upon heating, allowing more solute to dissolve.

Variables & Units:

• Fe = Iron (Solid Element)
• S = Sulphur (Solid Element)
• FeS = Iron(II) sulphide (Solid Compound)

Conditions/Notes:

• The product (FeS) is a compound with entirely different properties from its constituents.
• FeS is not attracted to a magnet, whereas elemental Iron was.
• Heating is strictly required (endothermic initiation) to force the chemical change.

Physical Changes:

• Occur without any alteration in the chemical composition or nature of the substance.
• No new substance is formed. The change is usually reversible.
• Examples: Melting ice, dissolving salt, boiling water.

Chemical Changes:

• Bring a change in the chemical properties, forming entirely new substances with different bonds.
• This is a chemical reaction. It is generally irreversible by physical methods.
• Examples: Rusting of iron, burning of wood, digestion.

Variables & Units:

• CaCO₃ = Calcium Carbonate (Solid)
• CaO = Calcium Oxide / Quicklime (Solid)
• CO₂ = Carbon Dioxide (Gas)

Conditions/Notes:

• This is a definitive chemical change.
• Calcium oxide (CaO) can be dissolved in water to form an alkaline/basic solution.
• Carbon dioxide (CO₂) dissolves in water to form a mild acidic solution.

Crystallization separates a pure solid in the form of its crystals from a solution.

It is superior to simple evaporation because:

1. Prevents Decomposition: Some solids (like sugar) may decompose or char when heated to complete dryness during evaporation.
2. Removes Soluble Impurities: Evaporation traps dissolved impurities within the solid mass. Crystallization leaves them behind in the residual liquid (mother liquor), yielding purely uniform crystals.

Principle: When a mixture is spun rapidly in a centrifuge, denser particles are forced to the bottom of the container, while lighter particles stay at the top due to centrifugal force.

Applications:

1. Separating cream from milk.
2. Medical: Diagnostic laboratories for blood and urine tests (separating cells from plasma).
3. Domestic/Industrial: Washing machines (the spin-drying cycle forces denser water outwards from clothes).

Principle: Used to separate solutes that dissolve in the same solvent. As the solvent rises up the stationary phase (filter paper) via capillary action, it carries solute particles. The component that is more soluble in the solvent travels faster and rises higher, thereby separating the mixture.

Applications:

1. Separating colours in a dye (e.g., black ink).
2. Extracting pigments from natural colours (e.g., chlorophyll).
3. Separating trace amounts of drugs from blood in forensic science.

The 25 K Rule:

• Simple Distillation: Used for miscible liquids that boil without decomposition and have a boiling point difference greater than 25 K.
• Fractional Distillation: Mandatory when the difference in boiling points is less than 25 K (e.g., separating components of air, or petroleum refining). It utilizes a fractionating column packed with glass beads to provide extra surface area for repeated condensation and vaporization.

Principle of Density Differences:
The separation relies strictly on the difference in densities of the two liquids.

When left undisturbed in a separating funnel, they separate into distinct layers. The denser liquid (e.g., water) settles at the bottom and can be carefully drained out via the stopcock, while the lighter liquid (e.g., oil) remains floating on top.

Sublimation is the direct transition from solid to gas without entering the liquid phase. It is used to separate a sublimable volatile component from a non-sublimable impurity.

Key Sublimable Solids:

1. Ammonium chloride ($NH_4Cl$)
2. Camphor
3. Naphthalene (mothballs)
4. Anthracene

Variables & Units:

• Volume of Solute = Volume of the dissolved liquid (mL or L)
• Volume of Solution = Total volume of solute + solvent (mL or L)

Conditions/Notes:

• Both volumes must be in the same SI or metric unit.
• Commonly used for liquid-liquid mixtures, such as alcohol in water.

The Tyndall effect is the scattering of a beam of light by colloidal particles (or very fine suspensions), which makes the actual path of the light beam distinctly visible.

Common Examples:

1. A fine beam of sunlight entering a dark room filled with suspended dust or smoke particles.
2. Sunlight passing through the canopy of a dense forest, where mist droplets act as colloidal particles dispersed in the air.

N/A - Merged

Mixture (Fe + S):

• Iron reacts independently with HCl to produce Hydrogen gas ($H_2$).
• Detection: Colourless, odourless, and burns with a distinctive "pop" sound. Sulphur remains unreacted.

Compound (FeS):

• The compound Iron(II) sulphide reacts with HCl to produce Hydrogen sulphide gas ($H_2S$).
• Detection: Colourless gas with a highly distinct smell of rotten eggs.

N/A - Merged

1. Sedimentation Tank: Water is allowed to stand still so that large, heavy suspended solid particles settle down under gravity.
2. Loading Tank: Chemicals (like alum) are added to coagulate finer, lighter suspended impurities, making them heavier so they can settle out.
3. Filtration Tank: The water passes through successive layers of fine sand, gravel, and coarse gravel to physically filter out any remaining micro-particles.
4. Chlorination: Chlorine is added to the filtered water to kill harmful bacteria and pathogens before it is supplied to homes.