Chapter 3.1 & 3.2 Notes

3.1 Exploring Diversity of Matter by its Chemical Composition (Elements, Compounds and Mixtures)

• Aims: classify substances as elements, compounds, and mixtures; define each; identify elemental building blocks; describe classification methods.

3.1.1 What is an element?

• Element: a pure substance that cannot be broken down into simpler substances by chemical methods.
• There are $118$ elements discovered; $92$ occur naturally on Earth; the rest are artificially made in laboratories.

3.1.2 How are elements represented?

• Symbols: usually the first one or two letters of the element; first letter capitalized, second letter lowercase.
• Why symbols? Each element has a unique symbol to avoid confusion.
• Examples: Hydrogen $ext{H}$, Carbon $ext{C}$, Chlorine $ext{Cl}$, Copper $ext{Cu}$, Iron $ext{Fe}$, Sodium $ext{Na}$.
• Notes:
  • Iron, copper and sodium symbols come from Latin names: ferrum, cuprum, natrium.
  • Chlorine is written as $ext{Cl}$ (capital C, lowercase l).

3.1.3 How can we classify and use elements?

• State of matter classification:
  • Gases: $ext{H}<em>2, ext{O}</em>2, ext{N}<em>2, ext{F}</em>2, ext{Cl}_2, ext{Ar}, ext{He}, ext{Ne}, ext{Kr}, ext{Xe}, ext{Rn}$
  • Liquids: $ext{Hg}, ext{Br}_2$
  • Solids: rest of the elements
• Metals vs non-metals (staircase line):
  • Metals: generally solid (except Hg), good conductors, malleable, shiny, hard, good heat conductors.
  • Non-metals: solids, liquids, or gases; poor conductors; brittle if solid; dull.
• Periodic Table organization:
  • Elements arranged by atomic number (proton number) in columns (groups) and rows (periods).
  • Similar properties in a group; properties change from metallic to non-metallic across a period.
  • Metals lie to the left of the staircase line; non-metals to the right.
  • Key symbols/terms: atomic number $Z$; elements grouped by similar properties.

3.1.4 Compounds

• Compound: a substance made up of two or more elements chemically joined.
• Examples and formulae (elements present):
  • Water: $ext{H}_2 ext{O}$ (hydrogen and oxygen)
  • Table salt: $ext{NaCl}$ (sodium and chlorine)
  • Nitrogen dioxide: $ext{NO}_2$ (nitrogen and oxygen)
  • Chalk (calcium carbonate): $ext{CaCO}_3$ (calcium, carbon, oxygen)
  • Sugar: $ext{C}<em>6 ext{H}</em>{12} ext{O}_6$ (carbon, hydrogen, oxygen)
  • Sand (silicon dioxide): $ext{SiO}_2$ (silicon and oxygen)

3.1.5 How are Compounds formed from Elements?

• Elements combine in fixed proportions to form compounds.
• Example: table salt formed from sodium and chlorine:
  • 1 ext{ g Na} + 1.54 ext{ g Cl}
    ightarrow 2.54 ext{ g NaCl}
  • 4 ext{ g Na} + 6.16 ext{ g Cl}
    ightarrow 10.16 ext{ g NaCl}
• Compounds have properties different from their constituent elements.
  • E.g., table salt is a white solid; sodium is a soft silvery metal; chlorine is a poisonous gas.

3.1.6 How can a Compound be broken down into simpler substances?

• Many compounds can be decomposed by heat or electricity.
• Products are simpler compounds or the constituent elements.
• Examples:
  • ext{CuCl}2 ightarrow ext{Cu} + ext{Cl}2
  • ext{CaCO}3 ightarrow ext{CaO} + ext{CO}2

3.1.7 Mixtures

• Mixture: two or more substances not chemically joined.
• Pure substances: contain only one element or one compound.
• All other substances are mixtures (can be elements, compounds, or both).
• Examples:
  • Bronze: copper + tin (a mixture of metals)
  • Table salt solution: table salt dissolved in water (a mixture of a compound in water)
  • Air: mixture of gases; some gases are elements (O2, N2), others are compounds (CO2, H2O).
• Properties of mixtures:
  • Substances in a mixture can be in any proportion.
  • Formation involves no chemical reaction (no heat/light exchange).
  • A mixture has properties of its components.
  • Can be separated by physical methods (filtration, distillation).

3.2 Solutions and Suspensions (Solutions, Suspensions and Solubility)

3.2.1(a) Solutions and Suspensions

• A mixture; two kinds:
  • Solutions
  • Suspensions
• Solutions: solute dissolved in solvent to form a homogeneous mixture.
• Suspensions: solid particles are not dissolved and remain suspended; heterogeneous.

3.2.1(b) How is a solution formed?

• Step 1: Solvent and solute come into contact.
• Step 2: Solute particles separate and mix evenly with solvent particles.
• Step 3: A homogeneous solution forms; solute particles are too small to be seen and pass through filter paper.
• Example: sugar in water forms a solution (sugar is soluble in water).

3.2.3 Solutes and Solvent

• Solvent: the liquid that dissolves the solute (e.g., water, alcohol).
• Solute: substance that dissolves in the solvent.
• In a sugar solution: Solute = $ext{C}<em>6 ext{H}</em>{12} ext{O}<em>6$; Solvent = $ext{H}</em>2 ext{O}$; Solution contains dissolved solute.

3.2.4 Properties of a solution

• Always clear; light passes through (tiny solute particles do not reflect light).
• May be coloured or colourless.
• Homogeneous: same colour, density, appearance throughout.
• If left to stand, no solid particles settle.
• Solute particles can pass through filter paper (too small to be trapped).

3.2.5(a) How is a suspension formed?

• When solids are mixed with water and do not dissolve, the mixture is a suspension (e.g., orange juice with pulp).

3.2.5(b) Properties of a suspension

• Usually cloudy; light cannot pass through.
• Heterogeneous: different properties in different parts.
• Solid particles settle to the bottom over time.
• Insoluble particles do not pass through filter paper.
• Note: solute particles in solutions are smaller than those in suspensions.

3.2.6 Comparison: Solution vs Suspension

• Solution: tiny solute particles dispersed evenly; homogeneous; light passes; no residue on filtration.
• Suspension: large insoluble particles; heterogeneous; cloudy; light blocked; particles may settle; filtration yields a residue.

3.2.7 Different types of solutions

• Dilute: small amount of solute in solvent.
• Concentrated: large amount of solute in solvent.
• Saturated: maximum amount of solute dissolved at a given temperature.

3.2.8 Formation of a saturated solution

• Method 1: add solute to solvent at a fixed temperature; stir until no more dissolves.
• Method 2: heat to evaporate solvent; crystallization on a cooled rod indicates saturation.
• Visual states:
  • Diluted solution
  • Concentrated solution
  • Saturated solution (excess solute not dissolved)

3.2.9 What are the factors affecting solubility?

• Solubility is the maximum mass of solute that can dissolve in a given volume of solvent at a given temperature.
• Factors:
  • Type of solute
  • Type of solvent
  • Temperature
• Examples:
  • Solubility of sucrose vs sodium chloride in water at a given temperature (sucrose more soluble than NaCl in the same conditions).
• Data (example at 70°C):
  • Mass of sugar dissolving in 100 g water: $325\ ext{g}$
  • Mass of sodium chloride dissolving in 100 g water: $30\ ext{g}$

3.2.10 What makes a substance dissolve faster?

• Rate of dissolving depends on:
  • Temperature: higher temperature speeds up dissolution.
  • Surface area: smaller pieces have greater surface area in contact with solvent; powders dissolve faster than chunks.
  • Stirring: stirring accelerates dissolution.

3.2.11 Applications of Solubility and Rate of Dissolving

• In tea, hot water plus stirring dissolves sugar quickly.
• In washing machines, soap and dirt dissolve faster with higher temperature and agitation; powders are used for faster dissolution.