AS Chem 1

Formulae, Equations and Amounts of Substance

  • The Mole

  • Key concept for chemical calculations.

  • Definition: Amount of substance in grams that equates to the number of particles found in 12 g of carbon-12.

  • Relative Atomic Mass (Ar)

  • Average mass of one atom, compared to 1/12 of the mass of carbon-12.

  • Molar Mass (Mr)

  • Mass in grams of one mole of a substance (g mol⁻¹).

  • Calculated by summing mass numbers from the periodic table.

  • Example: CaCO₃ = 40.1 (Ca) + 12.0 (C) + (16.0 x 3 (O)) = 100.1 g mol⁻¹.

  • Key Equations

  1. For pure solids, liquids, and gases: [ \text{amount} = \frac{\text{mass}}{\text{Molar Mass}} ]
  2. For solutions: [ \text{Concentration} = \frac{\text{amount}}{\text{volume}} ] (mol dm⁻³ or M)
  3. For gases: [ \text{Gas Volume (dm³)} = \text{amount} \times 24 ] (at 1 atm and 25°C)

  • Units

  • Mass: grams

  • Amount: mol

  • Concentration: mol dm⁻³ or M

  • Volume: dm³

  • Conversions

  • cm³ ➔ dm³: ÷ 1000

  • cm³ ➔ m³: ÷ 1,000,000

  • dm³ ➔ m³: ÷ 1000

  • Always express answers to 3 significant figures.

Example Calculations:

  • Calculate moles in 35.0 g of CuSO₄:

  • [ \text{Amount} = \frac{35}{(63.5 + 32 + (16 x 4))} \approx 0.219 \text{ mol} ]

  • Significant Figures: Answers should match the smallest significant figures of the data provided.

  • Changes of Units:

  • 1000 mg = 1 g

  • 1000 g = 1 kg

  • 1000 kg = 1 ton

  • Example Problem: Calculate moles in 75.0 mg of CaSO₄ . 2H₂O:

  • [ \text{Amount} = \frac{0.075}{(40 + 32 + (16 x 4) + (18 x 2))} \approx 4.36 x 10^{-4} \text{ mol} ]

Water of Crystallisation

  • Removal of Water: Can be achieved by heating (e.g., CaSO₄.xH₂O) which leads to dehydration.

  • Experiment Method:

    1. Weigh empty crucible.
    2. Add 2 g of hydrated CaSO₄.
    3. Heat and cool; Weigh until a constant mass is achieved.

  • Example: ZnSO₄.xH₂O, mass before and after gives ratio of water.

  • Example Calculation: Calculate integer x in ZnSO₄.xH₂O from moles calculated.

Avogadro's Constant

  • Definition: 1 mole = 6.02 x 10²³ particles (atoms, molecules, ions).

  • Example Calculation: Atoms in 6.00 g of tin:

  • [ \text{Amount} = \frac{6}{118.7} \approx 0.05055 \text{ mol} ]

  • Number of atoms: [ 0.05055 x 6.02 x 10^{23} \approx 3.04 x 10^{22} ]

  • Cl⁻ calculation in MgCl₂ solution:

  • Volume: [ 25.0 ext{ cm}^3 = 0.025 ext{ dm}^3 ]

  • Concentration: 0.400 mol dm⁻³, yielding 0.0200 mol of Cl⁻ ions.

Density Calculations

  • Density: g cm⁻³; mass can be calculated using [ \text{Mass} = \text{Density} \times ext{Volume} ].

  • Example Calculation: No. of ethanol molecules in 0.500 dm³ of ethanol:

  • Density: 0.789 g cm⁻³; mass = 394.5 g; amount = 8.576 mol.

Molecular and Empirical Formulas

  • Definitions:

  • Molecular Formula: Actual count of atoms in a compound.

  • Empirical Formula: Simplest ratio of atoms.

  • Molecular from Empirical:

  • Example: C₃H₆O (Mr = 116); C stands twice in Mr computation.

  • Molecular formula: C₆H₁₂O₂ from ratios.

  • Empirical Formula Calculation: K, I, O in a compound.

Concentration of Solutions

  • Equation: [ \text{Concentration} = \frac{\text{amount (mol)}}{\text{volume (dm}^3)} ].

  • Example Calculation: Solution made from 5.00 g Na₂CO₃ in 250 cm³: 0.189 mol dm⁻³.

  • Mass Concentration: g dm⁻³ = mol dm⁻³ × Mr.

Ions Dissociating in Solutions

  • Dissolution of solutes changes ion concentration.
  • Example: NaCl splits into Na⁺ and Cl⁻; MgCl₂ splits into Mg²⁺ and 2Cl⁻.

Making and Diluting a Solution

  • Dilution Steps: Measure, add water to a beaker and transfer.

  • Example: Calculate concentration after dilution; keep moles constant.

  • Final Calculations: Use original and new volume and concentration relationships.

Ideal Gas Equation: PV = nRT

  • Ideal for gas calculations; all gases follow at standard conditions.

  • Example Calculations: 100 kPa, 20°C, solving for mass of gases.

  • Understanding Gases: Real gas behavior vs ideal gas behavior.

Stoichiometry in Reactions

  • Using balanced equations to calculate quantities of substances.

  • Example Steps: Convert given data into mols, utilize stoichiometry to find unknowns.

Titrations and Calculations

  • Procedure: Clean equipment, measure accurately, observe color changes.

  • Calculations and Accuracy: Repeated titrations to find average, ensure correct reaction indicators are used.

Common Reactions and Equations

  • Neutralization, Acids and Bases, Precipitation reactions.

  • Example use: Common acid reactions and resulting products.

Hazards and Risks

  • Hazards: Definition and safety protocols in the lab.

  • Dealing with Excess Acids: Use of mild reagents to neutralize safely.