Gen chem 1 unit 7

General Information


Date: Monday, 07 Oct, 2024 (Sec. 01)Reference: OneNote Document (multiple pages)

Chemical Reactions

Electrolytes

Strong Electrolytes:

  • Definition: Strong electrolytes are substances that completely dissociate into ions when dissolved in water, leading to high conductivity.

  • Examples:

    • Strong acids: e.g., Hydrochloric acid (HCl), Hydrobromic acid (HBr) fully ionize to produce H⁺ and respective anions.

    • Strong bases: e.g., Group 1 metal hydroxides like Sodium hydroxide (NaOH), Potassium hydroxide (KOH) fully dissociate to generate OH⁻ ions in solution.

Weak Electrolytes:

  • Definition: Weak electrolytes partially dissociate in water, maintaining a considerable amount of the undissociated species in solution, and thus possess lower electrical conductivity.

  • Examples:

    • Weak acids: e.g., Acetic acid (HOAc), Hydrogen cyanide (HCN), Hydrofluoric acid (HF) only produce a fraction of ions in solution.

    • Weak bases: e.g., Ammonia (NH₃), Organic amine (OAE) produce few ions; however, they can accept protons from acids.

Non-Electrolytes:

  • Definition: Non-electrolytes are molecular compounds that do not form ions in solution, resulting in no conduction of electricity.

  • Examples:

    • Substances: e.g., Water (H₂O), Carbon dioxide (CO₂), Urea do not dissociate into ions when added to water.

Types of Reactions

  • Driving Forces in Reactions:

    • Chemical reactions in aqueous solutions typically involve ions.

    • Reactive Ions: Ions that actively participate in a reaction, undergoing changes in charge.

    • Spectator Ions: Ions present in the solution that do not undergo any change during the chemical reaction and can be removed from the net ionic equation.

  • Ways to Express Chemical Reactions:

    • Molecular Reaction: Displays all reactants and products in the chemical reaction, including molecules and ions.

    • Full Ionic Reaction: Illustrates all ions present in the reaction; crucial for understanding all charged species involved.

    • Net Ionic Reaction: Simplifies the reaction to show only the ions and molecules that undergo a chemical change.

Precipitation and Acid/Base Reactions

  • Precipitation Reactions:

    • Involves the formation of an insoluble solid, known as a precipitate, from reactants dissolved in solution.

    • Example Reaction:

      • Molecular: 2KI(aq) + Pb(NO₃)₂(aq) → PbI₂(s) + 2K⁺(aq)

      • Full Ionic: 2K⁺(aq) + 2I⁻(aq) + Pb²⁺(aq) → PbI₂(s) + 2K⁺(aq)

      • Net Ionic: 2I⁻(aq) + Pb²⁺(aq) → PbI₂(s)

Titration

Overview

  • Titration: A quantitative analytical method used to determine the concentration of a solution by reacting it with a standard solution of known concentration.

  • End Point: The stage in titration where the reaction has reached completion, often indicated by a color change due to the pH indicator used.

Acid/Base Titration Examples

  • Example 1: Neutralization of hydrochloric acid (HCl) with sodium hydroxide (NaOH).

    • Given: 0.125 M NaOH needed to neutralize 25.0 mL of 0.750 M HCl.

  • Calculation:

    • Moles of HCl = 0.750 mol/L * 0.025L = 0.01875 mol.

    • Required NaOH = 0.01875 mol (based on 1:1 stoichiometry of neutralization).

    • Volume of NaOH = 0.01875 mol / 0.125 mol/L = 0.15L = 150 mL.

Indicators in Titrations

  • Various indicators are used in titrations, each changing color at specific pH levels, ensuring accurate detection of the end point.

    • Examples: Phenolphthalein turns pink in basic solutions, while Methyl Orange turns red in acidic solutions and yellow in neutral to basic solutions.

Redox Reactions

Overview

  • Redox Reactions: Reactions involving the transfer of electrons between chemical species, essential in both organic and inorganic chemistry.

  • Oxidation: The process involving the loss of electrons, leading to an increase in oxidation state.

  • Reduction: The gain of electrons, resulting in a decrease in oxidation state.

  • The driving force of redox reactions is the decrease in electrical potential energy.

Oxidation States

  • Oxidation Number Assignment Rules:

    • Pure elements have an oxidation number of 0 (e.g., O₂, N₂).

    • Simple ions possess oxidation numbers equal to their charges (e.g., Na⁺ = +1, Cl⁻ = -1).

    • In compounds, the more electronegative atom receives a negative oxidation state based on its position in the periodic table.

    • The sum of oxidation numbers in a neutral compound is always zero.

    • The total oxidation number in a polyatomic ion equals the charge of that ion.

Balancing Redox Reactions

  • Half-Reaction Method:

    • Start by splitting the overall redox reaction into oxidation and reduction half-reactions.

    • Balance all atoms other than hydrogen (H) and oxygen (O) first.

    • Balance oxygen by adding water (H₂O) molecules, and balance hydrogen by adding hydrogen ions (H⁺).

    • Balance the charge on both sides using electrons.

    • Combine the half-reactions and simplify where possible.

  • Example Balancing in Acidic Solutions:

    • Reaction: MnO₄⁻(aq) + I₂(aq) → Mn²⁺(aq) + IO₃⁻(aq)

    • Steps:

      • Balance Manganese (Mn), Iodine (I), Hydrogen (H), and Oxygen (O) as necessary.

      • Determine oxidation states for each species and balance charges as needed.