(481) Unit 4.2 - Net-Ionic Equations

Introduction to Net Ionic Equations

  • Focus of Unit 4.2: Writing net ionic equations.

  • Importance of having periodic tables handy.

Key Principles of Chemical Reactions

  • Balancing Reactions: Essential for chemical reactions to represent the conservation of mass.

    • If coefficients are present, reactions are likely balanced.

    • If no coefficients, verify the balance.

  • Types of Reactions: Differentiate between chemical and physical reactions.

    • Reactions typically involve changes in chemical substances, but physical processes like dissolution can also be represented.

Example of Dissolution

  • Calcium Hydroxide (Ca(OH)2):

    • Question: Write a balanced equation for its dissolution in water.

    • Understanding Dissolution: The process is physical, not chemical; water does not participate in the reaction directly.

    • Balanced dissolution reaction:

      • Ca(OH)2 (s) → Ca²⁺ (aq) + 2 OH⁻ (aq)

    • States of matter notation included for clarity, but not mandatory for AP exam.

Particle Representation

  • Particle Representation: Diagram with water molecules around ionized calcium.

    • Water is a polar molecule, with the negatively charged oxygen oriented towards positive ions.

    • Importance of retaining memory regarding molecular orientation from previous units.

Introduction to Net Ionic Equations

  • Molecular Equation Example: Potassium iodide (KI) + Lead nitrate (Pb(NO3)2) → Lead iodide (PbI2) + Potassium nitrate (KNO3)

  • States: Aqueous (dissolved) and solid (precipitate).

    • KI, Pb(NO3)2, and KNO3 are soluble (dissolve), but PbI2 is insoluble (precipitate).

  • Complete Ionic Equation: Separate soluble substances into their respective ions.

    • Example:

      • 2K⁺ (aq) + 2I⁻ (aq) + Pb²⁺ (aq) + 2NO3⁻ (aq) → PbI2 (s) + 2K⁺ (aq) + 2NO3⁻ (aq)

  • Spectator Ions: Identify ions that do not change during the reaction—cancel them out.

    • Example: K⁺ and NO3⁻ are spectator ions.

  • Net Ionic Equation: The final equation showing only the ions involved in forming the precipitate:

    • Pb²⁺ (aq) + 2I⁻ (aq) → PbI2 (s)

Solubility Rules

  • Understanding Solubility: Importance for predicting outcomes in reactions.

    • General Rules: Several compounds are generally soluble:

      • a) Nitrate (NO3⁻) and acetate (C2H3O2⁻) compounds are always soluble.

      • b) Salts of alkali metals and ammonium (NH4⁺) are also soluble.

  • Exceptions: Some ions have exceptions; it’s not enough to rely solely on general rules.

    • For example: Most sulfate (SO4²⁻) salts are soluble except for those with Ba²⁺, Sr²⁺, and Pb²⁺ ions.

  • Hydroxides and Others:

    • Most hydroxide salts are insoluble except for those containing alkali metals.

  • Strong Acids: Memorize six strong acids that ionize completely in solution.

Conflicting Information in Solubility

  • Example: Sodium sulfide (Na2S) is soluble because the rule regarding alkali metals supersedes the rule about sulfide salts being mostly insoluble.

Practice Problem**

  • Task: Write three types of equations for iron (III) sulfate and potassium sulfide.

    • Predict Products: Iron (III) sulfate and potassium sulfide → potassium sulfate + iron (III) sulfide.

  • Consider solubility of sodium sulfide and its potential conflict.

  • Single Replacement Reaction: Explanation of how to write equations involving diatomic molecules.

  • Final Notes: Understanding that Cl2 cannot be simplified to Cl-; those are fundamentally different.

    • Net ionic equation conclusion: Identify and retain key ions after canceling out spectators.