Ch - 13 Revision ppt

Chapter 13: Compounds in Aqueous Solution

Overview of the chapter’s focus on aqueous solutions, including dissociation and ionization of compounds.
Initial comparison of solid crystal particles of CuSO4•5H2O versus those in solution for better understanding.

Write dissolution equations for soluble ionic compounds in water.
Predict precipitate formation when mixing soluble ionic compounds, and write net ionic equations for precipitation reactions.
Distinguish between dissociation of ionic compounds and ionization of molecular compounds.
Draw and explain the hydronium ion structure and its significance.
Define and differentiate between strong and weak electrolytes.

Definition: Dissociation is the separation of ions that occurs when an ionic compound dissolves in water.
Example: Dissociation of NaCl into Na⁺ and Cl⁻ ions in solution, with the equation representing the process.

Sample Problem A: Dissolution of Al2(SO4)3 in water.
- 1 mol of aluminum sulfate produces 2 moles of aluminum ions and 3 moles of sulfate ions.
- Total of 5 moles of ions produced from 1 mol of aluminum sulfate.

Concept: No compound is completely insoluble, but very low-solubility compounds can be treated as such.
General Solubility Guidelines:
1. Soluble: Sodium, potassium, and ammonium compounds.
2. Nitrates, acetates, and chlorates are soluble.
3. Most chlorides are soluble except those of Ag, Hg(I), and Pb.
4. Most sulfates are soluble except those of Ba, Ca, Hg, Sr, and Pb.
5. Most carbonates and phosphates are insoluble except for those of Na, K, and NH4.
6. Insoluble sulfides except for those of Ca, Sr, Na, K, and NH4.

Listed examples of soluble (e.g., NiCl2, KMnO4) versus insoluble compounds (e.g., AgCl, CdS).

Definition: Net ionic equations reflect only the ions and compounds that change during a reaction, ignoring spectator ions.
- Importance in simplifying reactions involving aqueous solutions.
Example equations provided to illustrate how to write these based on given compounds and reactions.

Definition: Ionization is a process where solute molecules form ions in solution due to solvent action, creating ions from molecular compounds that did not previously exist.
Example: Hydrogen chloride (HCl) ionizing in aqueous solution. Presence of hydronium ion (H3O+).

Electrolytes: Substances yielding ions in solution and conducting electricity.
Strong Electrolytes: Dissociate completely in water, yielding ions and conducting well (e.g., HCl, HBr, HI).
Weak Electrolytes: Partially dissociate, resulting in fewer ions in solution (e.g., HF).

Definition: Properties depend on solute particle concentration, regardless of particle identity, include:
- Vapor-Pressure Lowering
- Freezing-Point Depression
- Boiling-Point Elevation
- Osmotic Pressure

Nonvolatile solutes reduce vapor pressure, affecting boiling and freezing points of solutions.

The freezing-point difference between pure solvent and solution is proportional to molal concentration: ∆tf = Kfm.
Example: Calculation of freezing-point depression applied to given problems.

Similar to freezing-point depression: Directly relates the concentration of solute to the boiling point elevation with ∆tb = Kbm.

The concept of semipermeable membranes in controlling solute passage; defined as pressure needed to halt osmosis, with practical applications.

Electrolytes generate multiple particles per mole of compound in solution, affecting colligative properties beyond expectations.
Comparison example between nonelectrolytes (like sucrose) and electrolytes (like NaCl).
Influences of attractive forces between ions leading to deviations from expected properties.

A series of sample problems provided to demonstrate practical applications of concepts, including freezing-point depression, boiling-point elevation, and calculations involving molality.

Key takeaways include the importance of understanding dissociation, ionization, and how colligative properties are influenced by the solute nature in solutions.