CHEMISTRY LECTURE

JAMB CHEMISTRY LECTURE

1. Separation of Mixtures and Purification

(a) Pure and Impure Substances

A pure substance contains only one type of particle (element or compound), while an impure substance (mixture) contains two or more substances physically combined.

• Pure substances have fixed melting and boiling points.
• Impure substances melt over a range.

(b) Boiling and Melting Points

Pure substances have sharp melting points. Impurities lower melting point and increase boiling range.

(c) Elements, Compounds, Mixtures

• Element: Cannot be broken chemically (e.g., O₂)
• Compound: Chemically combined (e.g., H₂O)
• Mixture: Physically combined (e.g., sand + salt)

(d) Physical vs Chemical Changes

• Physical: No new substance (melting ice)
• Chemical: New substance formed (burning)

(e) Separation Techniques

• Evaporation – recovering solute
• Distillation – separating liquids
• Fractional Distillation – different boiling points
• Filtration – solid from liquid
• Chromatography – separates based on movement
• Sublimation – solid → gas (e.g., iodine)
• Decantation – pouring off liquid

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2. Chemical Combination (Stoichiometry)

Laws

• Law of Conservation: Mass is neither created nor destroyed
• Definite Proportion: Same ratio always
• Multiple Proportion: Simple ratios
• Gay Lussac: Volumes combine in ratios
• Avogadro: Equal volumes → equal molecules

Mole Concept

1 mole = 6.02 × 10²³ particles

Example 1

Calculate moles in 44g of CO₂

Molar mass CO₂ = 44 Moles = 44 / 44 = 1 mole

Example 2

Calculate number of molecules in 2 moles of O₂

= 2 × 6.02 × 10²³ = 1.204 × 10²⁴ molecules

Stoichiometric Calculation

2H₂ + O₂ → 2H₂O

If 4 moles of H₂ react: Ratio H₂:H₂O = 2:2 → 4 moles produce 4 moles H₂O

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3. Concentration Calculations

g/dm³

Mass concentration = mass / volume

Example: 10g in 2dm³ = 10 / 2 = 5 g/dm³

mol/dm³

Molar concentration = moles / volume

Example: 0.5 moles in 0.25dm³ = 0.5 / 0.25 = 2 mol/dm³

Conversion (25cm³ → 1000cm³)

Factor = 1000 / 25 = 40 Multiply concentration by 40

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4. Kinetic Theory of Matter

• Particles are always moving
• Higher temperature → faster motion

States Explained

• Melting: solid → liquid
• Boiling: liquid → gas
• Freezing: liquid → solid
• Condensation: gas → liquid

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5. Gas Laws

Boyle's Law

P₁V₁ = P₂V₂

Example: P₁ = 2 atm, V₁ = 4 dm³ P₂ = ?, V₂ = 2 dm³

2 × 4 = P₂ × 2 P₂ = 4 atm

Charles Law

V₁/T₁ = V₂/T₂

Ideal Gas Equation

PV = nRT

Example: P = 1 atm, V = 22.4 dm³, R = 0.0821, T = 273K

n = PV / RT = (1 × 22.4) / (0.0821 × 273) = 1 mole

Vapour Density

VD = Molar Mass / 2

Example: VD = 14 Molar mass = 2 × 14 = 28

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EXAM SUMMARY

• Use mole = mass / molar mass
• Use PV = nRT for gases
• Always balance equations
• Convert cm³ to dm³ (÷1000)
• Use correct separation method