Q2_L3_ Stoichiometry_GEN CHEM 1

Stoichiometry

• Definition: The mathematical relationship between the quantities of reactants and products in a chemical reaction.

• Key Components: Masses, moles, and percentages can be calculated within a chemical equation.

• Importance: Essential for understanding the proportions of substances in reactions.

Formula and Molecular Mass

• Formula Mass (Formula Weight):

  • Sum of atomic weights of atoms in a compound's empirical formula.

• Molecular Mass (Molecular Weight):

  • Average mass of a molecule calculated from its molecular formula.

• Example Calculation:

  • For calcium chloride (CaCl2):

    • Ca: 1(40.1 amu) + Cl: 2(35.5 amu) = 111.1 amu

Percentage Composition

• Definition: Measurement of each element's mass percentage in a compound.

• Calculation Formula:

  • % element = (number of atoms) x (atomic weight) x 100 / (formula weight of the compound)

• Example:

  • In ethane (C2H6):

    • %C = (24.0 amu / 30.0 amu) x 100 = 80.0%

    • %H = (6.06 amu / 30.0 amu) x 100 = 20.0%

Types of Formulas

1. Empirical Formula:

• Simplest ratio of atoms in a compound (e.g., NaCl).

2. Molecular Formula:

• Total number of each atom in a covalent compound (e.g., C2H6).

3. Structural Formula:

• Arrangement of atoms in a compound (e.g., CH4).

Deriving Empirical and Molecular Formulas

• Empirical Formula:

  • Based on mass percentages of elements. Steps:

    1. Assume a 100g sample to convert percentages to grams.

    2. Convert grams to moles using atomic masses.

    3. Divide by the smallest mole fraction to find the simplest ratio.

• Example:

  • For a sample containing 43.4% Na, 11.3% C, and 45.3% O:

    • Na: 1.887 moles, C: 0.942 moles, O: 2.831 moles

    • Empirical Formula = Na2CO3.

Additional Example for Empirical Formula

• For Calcium:

  • Example of 8.00 g Ca + 3.20 g O gives empirical formula CaO.

Calculation Method for Molecular Formula

• Using Empirical Formula:

  • Example: For propylene with 14.3% H and 85.7% C, with molar mass = 42.0 g/mol:

    • Empirical Formula = CH2;

    • Multiplier = 42 g/mol / 14 g/mol = 3.

    • Molecular Formula = C3H6.

The Mole Concept

• Definition: A unit for counting particles.

  • 1 mole = Avogadro’s Number (6.02 x 10²³).

• Molar Mass: Mass of 1 mole of a substance (g/mol).

• Mole Relationships:

  • Mass per mole is equal to its formula/molecular mass.

Chemical Reactions

Five Types of Chemical Reactions

1. Combustion: Involves burning in oxygen.

2. Synthesis: Combination of substances to form a complex product (A + B = AB).

3. Decomposition: Breakdown of a compound into simpler substances (AB = A + B).

4. Single Displacement: An element in a compound is replaced (A + BC = AC + B).

5. Double Displacement: Exchange of ions between two compounds (AB + CD = AD + CB).

Evidence of a Chemical Reaction

• Color/odour change, gas/solid formation, difficult to reverse, energy release/absorption.

Reactants in Solutions

• Solute: The substance being dissolved.

• Solvent: The substance doing the dissolving (usually water).

• Aqueous Solution: A solution where water is the solvent.

• Solubility: Measure of how much solute can dissolve in a solvent.

Solubility Table Usage

• Determines solubility of various compounds based on ion composition.

Common Alloy Percent Composition

• Brass: 70% Cu, 30% Zn - Durable and corrosion-resistant.

• Stainless Steel: 80% Fe, 18% Cr, 1% Ni, 1% Si – Corrosion-resistant.

• 18K Gold: 75% Au, 13% Ag, 12% Cu – Harder than 24K gold.

Ionic Compounds

• Formed by metals donating electrons and nonmetals accepting electrons, creating cations and anions (e.g., NaCl).

• Writing chemical formulas involves identifying the charge balance between ions.