20240410_Reactions(1 per page Colour)

Chemical Reactions Overview

• Conceptual Framework:

  • A re-distribution of atoms occurs during chemical reactions.

  • Existing chemical bonds break, and new bonds form.

  • Based on conservation laws:

    • Conservation of mass.

    • Conservation of atoms.

• Foundational Contributors:

  • Antoine Lavoisier (1743 - 1794) emphasized these conservation principles.

Writing and Balancing Equations

• Balanced Chemical Reactions:

  • Can be expressed in equations, e.g., 4P + 6Cl2 -> 4PCl3.

  • Must balance for both mass and charge.

  • Example:

    • H3PO2(aq) + 2H2O(ℓ) → H3PO4(aq) + 4H+(aq)

    • (Balanced regarding mass and charge)

• Stoichiometry and Coefficients:

  • Relative amounts of reactants and products are expressed using stoichiometric coefficients.

  • From a balanced equation:

    • Theoretical calculations can be made on reactions.

Limitations of Balanced Equations

• Do Not Provide:

  • Actual amounts of reactants and products.

  • Natural tendency of a reaction to occur.

  • Energy changes associated with the reaction.

  • Rate of reaction or its mechanism.

Reactions in Solution

• Aqueous Solutions:

  • Homogeneous mixtures of solutes dissolved in solvents (typically water).

  • Essential for various geological, geographical, and biological processes.

• Water as a Solvent:

  • Inexpensive and capable of dissolving many substances.

  • Aids in the dissociation of solutes (e.g., NaCl splits into Na+ and Cl- ions).

Electrolytes in Solution

• Types of Electrolytes:

  1. Strong Electrolytes: Completely ionized, good conductors of electricity (e.g., NaCl).

  2. Non-electrolytes: Do not form ions in solution, poor electrical conductors (e.g., sucrose).

  3. Weak Electrolytes: Partially ionized in solution, with limited electrical conductivity (e.g., acetic acid).

Solubility Rules and Precipitation Reactions

• Overview:

  • Many ionic compounds are not soluble in water.

  • Certain combinations of cations and anions lead to precipitation reactions, forming insoluble compounds (e.g.,

    • Formation of insoluble salts, e.g., BaSO4 from mixtures).

• Equations in Precipitation Reactions:

  • Use molecular equations, full ionic equations, and net ionic equations to express reactions.

  • Example:

    • Ca2+(aq) + SO4^2−(aq) → CaSO4(s), where CaSO4 precipitates.

Acid-Base Reactions

• General Concepts:

  • Acids produce H+ in solution, while bases produce OH-.

  • Importance in home, biology, industry, and environmental processes.

• Definitions:

  • Arrhenius Model:

    • Acids increase H+ concentration in water (e.g., HCl).

    • Bases increase OH- concentration in water (e.g., NaOH).

• Brønsted-Lowry Model:

  • Defines acids as proton donors and bases as proton acceptors.

Redox Reactions

• Basic Principles:

  • Coupling of oxidation (gain of oxygen/loss of electrons) and reduction (loss of oxygen/gain of electrons).

  • Recognize oxidation states for reactions:

    • In metal extraction processes (e.g., extraction of iron), recognize reducing and oxidizing agents.

• Balancing Redox Reactions:

  • Involves assigning oxidation states, splitting into half-reactions, and balancing mass and charge separately.

• Examples of Balancing:

  • Example: Aℓ + Ni2+ → Aℓ3+ + Ni:

    • Assign oxidation states to identify oxidation and reduction.

Conclusion

• Understanding Chemical Reactions:

  • Concepts of conservation, balancing equations, the role of solvents, electrolytes, and definitions of acids and bases are fundamental.

  • Focus on redox reactions and how they apply to practical scenarios in both laboratory and real-world chemistry.