Chemistry Chapter: Moles and Stoichiometry

Moles of Solution

• Moles of solution calculated using the formula: Moles = Concentration × Volume.
• Concentration units: mol/dm³ or g/dm³; Volume unit: dm³.

Empirical & Molecular Formula

• Empirical Formula: Simplest ratio of atoms of each element in a compound.
• Example: Copper oxide with 8g Cu and 1g O.
• Molecular Formula: Actual number of atoms of each element in a molecule.
• Relative Molecular Mass (Mr) calculated as: n × Relative Mass of Empirical Formula.
• Example: For propene (CH₂) with Mr 42, n calculated as 3.

Stoichiometry

• Emphasis on mole ratios; answers should be in decimals with appropriate significant figures.
• Example of stoichiometry problem with CaO and water leading to Ca(OH)₂, calculating maximum mass using mole ratio.
• Mass conversions: 1 tonne = 1000 kg; 28 tonnes = 28,000 kg = 28,000,000 g.
• Balanced equation required for accurate calculations.

Limiting & Excess Reactants

• Many reactions use excess amounts of one reactant.
• Example of copper purity calculation based on mass of CuO present; % purity found to be 84%.

Percentage Yield

• Theoretical yield is mass calculated from reaction equations; actual yield is obtained from experiments.
• Actual mass is often lower than theoretical due to issues like incomplete reactions or losses during the process.
• Example calculation of percentage yield based on 28g of CuSO₄ from 0.2 mol of CuO, resulting in a yield of 87.5%.

Acids & Bases

• Acids: Produce H⁺ ions in water, sour taste, conduct electricity, pH < 7, and turn blue litmus red.
• React with metals, bases, and carbonates:
  • Acid + Reactive Metal → Salt + Hydrogen.
  • Acid + Base → Salt + Water.
  • Acid + Metal Carbonate → Salt + Water + CO₂.
• Solid acid salts release H⁺ ions in water but behave differently in organic solvents.

Salts

• Ionic compounds made of cations and anions formed by replacing H⁺ ions in acids.
• Preparation methods vary based on solubility of salt and reagents:
  • Precipitation for insoluble salts.
  • Acid + Excess Solid for soluble salts containing Group 1/Ammonium.
  • Titration for other soluble salts.

Precipitation Method

• Mix two soluble solutions to form an insoluble salt.
• Standard procedure: Determine solutions, mix, filter, rinse, and dry precipitate.

Acid + Excess Solid Method

• For soluble salts not containing Group 1/Ammonium ions.
• Solid must be in excess to ensure complete utilization of acid.

Oxidation States

• Group 1 metals: +1 oxidation state.
• Group 2 metals: +2 oxidation state.
• Specific non-metals have defined states: Fluorine (-1), Hydrogen (+1 or -1), Oxygen usually (-2).
• Transition metals and lanthanides typically have positive oxidation states.

Identifying Reactions

• Redox Reactions involve oxidation and reduction (e.g., respiration, photosynthesis, and the Haber process).
• Non-redox reactions, such as acid-base reactions (neutralization), show no change in oxidation states.

Endothermic and Exothermic Reactions

• Endothermic: Energy absorbed, temperature decreases, ∆H > 0.
• Exothermic: Energy released, temperature increases, ∆H < 0.

Bond Energy:

• Amount of energy related to breaking or forming chemical bonds.
• Energy change calculated as: ∆H = Total Energy absorbed - Total Energy released.

Periodic Table Overview

• Elements arranged by increasing atomic number in periods (horizontal) and groups (vertical).
• Groups share similar valence electrons and chemical properties.
• Across a period: decrease in metallic properties, increase in non-metallic properties.
• Down a group: increase in metallic properties, decrease in non-metallic properties.

Test your knowledge on calculating moles of a solution, determining empirical and molecular formulas, and understanding stoichiometry. This quiz also covers limiting and excess reactants, providing fundamental concepts essential for mastering chemical reactions. Prepare to apply these concepts to solve real-world chemical problems.