Solubility Rules and Ionic Compounds

Soluble and Insoluble Ionic Compounds

Understanding solubility rules is essential for predicting the outcomes of precipitation reactions and for writing net ionic equations.

• Soluble Ionic Compounds:

  • All common compounds of Group 1A ions (alkali metals) and ammonium ion (NH4+) are soluble.

  • All common nitrates (NO3-), acetates (CH3COO-), and most perchlorates (ClO4-) are soluble.

  • All common chlorides, bromides, and iodides are soluble except those of Ag+, Pb2+, Cu+, Hg22+.

  • All common fluorides are soluble except those of Pb2+ and Group 2A elements.

  • All common sulfates (SO42-) are soluble except those of Ca2+, Sr2+, Ba2+, Ag+, Pb2+.

• Insoluble Ionic Compounds:

  • All common metal hydroxides are insoluble except those of Group 1A and the larger members of Group 2A (beginning with Ca2+).

  • All common carbonates (CO32-) and phosphates (PO43-) are insoluble except those of Group 1A and ammonium.

  • All common sulfides (S2-) are insoluble except those of Group 1A, Group 2A, and ammonium.

Application: Use these rules to predict whether a precipitate will form in double displacement reactions.

Electrolytes and Acid-Base Classification

Strong, Weak, and Non-Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water due to the presence of ions.

• Strong Electrolytes: Completely dissociate into ions (e.g., NaCl, HCl, NaOH).

• Weak Electrolytes: Partially dissociate (e.g., NH3, HF, CH3COOH).

• Non-Electrolytes: Do not dissociate (e.g., CH3OH, sugar).

Examples: Classify the following:

• NH3: Weak base, weak electrolyte

• NaOH: Strong base, strong electrolyte

• NaCl: Strong electrolyte (ionic salt)

• HCl: Strong acid, strong electrolyte

• H2SO4: Strong acid, strong electrolyte

• H3PO4: Weak acid, weak electrolyte

• HF: Weak acid, weak electrolyte

• HBr: Strong acid, strong electrolyte

• H2CO3: Weak acid, weak electrolyte

• CH3OH: Non-electrolyte

• CH3COOH: Weak acid, weak electrolyte

Acids and Bases

• Acid: Substance that donates a proton (H+).

• Base: Substance that accepts a proton or donates OH-.

• Strong Acids (memorize): HCl, HBr, HI, HNO3, HClO4, H2SO4, HClO3

• Strong Bases: Group 1A and heavy Group 2A hydroxides (e.g., NaOH, KOH, Ca(OH)2).

Writing Chemical Equations

Types of Equations

• Molecular Equation: Shows all reactants and products as compounds.

• Total Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the species that change during the reaction.

Note: Weak acids and bases are not split into ions in ionic equations.

Neutralization and Gas Formation

• Neutralization: Acid + Base → Salt + Water

• Gas Formation: Acids react with carbonates/bicarbonates to produce CO2 gas.

Example: HCl + NaHCO3 → NaCl + H2O + CO2

Molarity, Dilutions, and Solution Stoichiometry

Molarity and Dilution Calculations

• Molarity (M):

• Dilution Equation:

Application: Use these equations to prepare solutions of desired concentrations.

Stoichiometry in Solution

• Use molarity to convert between volume and moles in reactions.

• Apply stoichiometric coefficients from balanced equations.

Titration Concepts

• End Point: The point at which the indicator changes color.

• Equivalence Point: The point at which stoichiometrically equivalent quantities of acid and base have reacted.

• Indicator: A substance that changes color at (or near) the equivalence point.

Example: Given titration data, calculate the unknown concentration using stoichiometry.

Acid-Base Equilibria (Chapter 16)

Acid-Base Definitions

• Arrhenius Acid: Produces H+ in water.

• Arrhenius Base: Produces OH- in water.

• Brønsted-Lowry Acid: Proton donor.

• Brønsted-Lowry Base: Proton acceptor.

• Conjugate Acid-Base Pairs: Differ by one proton.

Acid and Base Strength

• The stronger the acid, the weaker its conjugate base, and vice versa.

Autoionization of Water and Kw

• Water self-ionizes:

• Ion-product constant: at 25°C

pH and pOH Calculations

Calculating pH for Strong and Weak Acids/Bases

• For strong acids/bases, assume complete dissociation.

• For weak acids/bases, use ICE tables and equilibrium constants (, ).

• Percent ionization:

Relationship Between Ka, Kb, and Kw

pKa and Acid Strength

• Lower pKa = stronger acid.

• pKa is related to Ka:

Acid-Base Properties of Salts

• Salts can be acidic, basic, or neutral depending on the strengths of their constituent acids and bases.

Predicting Acid Strength from Structure

• Factors: Bond strength, electronegativity, resonance stabilization, and inductive effects.

Buffers and Titrations (Chapter 17)

Buffer Solutions

Buffers are solutions that resist changes in pH upon addition of small amounts of acid or base.

• Buffer Capacity: The amount of acid or base a buffer can neutralize before pH changes significantly.

• Buffer Range: The pH range over which the buffer is effective (typically pKa ± 1).

Henderson-Hasselbalch Equation

Use this equation to calculate buffer pH and to determine the ratio of [A-] to [HA] for a desired pH.

Buffer Calculations with Added Acid or Base

• Use ICE tables to account for the reaction of added strong acid/base with buffer components.

Titration Curves and pKa Determination

• At the half-equivalence point: (when )

• For polyprotic acids, each equivalence point corresponds to the loss of one proton.

Types of Titration Curves

• Weak acid–strong base

• Polyprotic acid–strong base

• Weak base–strong acid

Indicators and Equivalence Point Selection

• Select an indicator whose color change range includes the equivalence point pH.

Summary Table: Solubility Rules

Additional info: ICE tables (Initial, Change, Equilibrium) are used to solve equilibrium problems for weak acids and bases. For titrations, the equivalence point is where moles of acid equal moles of base. Buffer solutions are most effective when the concentrations of acid and conjugate base are similar, and the pH is close to the pKa of the acid.