Ch. 4 Aqueous Solutions, Electrolytes, Acids & Bases, and Stoichiometry: General Chemistry Study Notes

General Properties of Aqueous Solutions

Solute, Solvent, and Solution

An aqueous solution is a homogeneous mixture in which water acts as the solvent. The solute is the substance dissolved in the solvent.

• Solute: Substance being dissolved (e.g., NaCl).

• Solvent: Substance doing the dissolving (e.g., H2O).

• Solution: Homogeneous mixture of solute and solvent.

• Example: Table salt (NaCl) dissolved in water forms an aqueous solution.

Processes When Ionic Compounds Dissolve

When ionic compounds dissolve in water, the ions separate and become surrounded by water molecules (hydration).

• Dissociation: Ionic solids break into ions (e.g., NaCl → Na+ + Cl-).

• Hydration: Water molecules surround and stabilize ions.

• Example: NaCl crystal dissolving in water forms hydrated Na+ and Cl- ions.

Electrolytes in Aqueous Solutions

Strong Electrolytes

Strong electrolytes dissociate completely in water, producing solutions that conduct electricity well.

• Strong Acids: HCl, HNO3, H2SO4

• Strong Bases: NaOH, KOH

• Most soluble ionic salts

• Equation: $\text{NaCl} (s) \rightarrow \text{Na}^+ (aq) + \text{Cl}^- (aq)$

Weak Electrolytes

Weak electrolytes only partially dissociate in water, resulting in poor conductivity.

• Weak Acids: CH3COOH (acetic acid), HF

• Weak Bases: NH3 (ammonia)

• Equation: $\text{CH}_3\text{COOH} (aq) \rightleftharpoons \text{CH}_3\text{COO}^- (aq) + \text{H}^+ (aq)$

Nonelectrolytes

Nonelectrolytes dissolve in water but do not produce ions, so their solutions do not conduct electricity.

• Examples: Sugar (C12H22O11), ethanol (C2H5OH)

• Equation: $\text{C}_6\text{H}_{12}\text{O}_6 (s) \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 (aq)$

Acids and Bases in Aqueous Solutions

Arrhenius and Brønsted-Lowry Definitions

Acids and bases are defined by their behavior in water.

• Arrhenius Acid: Produces H+ ions in water.

• Arrhenius Base: Produces OH- ions in water.

• Brønsted-Lowry Acid: Donates a proton (H+).

• Brønsted-Lowry Base: Accepts a proton (H+).

Strong vs. Weak Acids and Bases

Strong acids and bases dissociate completely, while weak acids and bases only partially dissociate.

• Strong Acids: HCl, HNO3, H2SO4

• Weak Acids: CH3COOH, HF

• Strong Bases: NaOH, KOH

• Weak Bases: NH3

• Equation (Strong Acid): $\text{HCl} (aq) \rightarrow \text{H}^+ (aq) + \text{Cl}^- (aq)$

• Equation (Weak Acid): $\text{HF} (aq) \rightleftharpoons \text{H}^+ (aq) + \text{F}^- (aq)$

Neutralization Reactions

Acids and bases react to form water and a salt in a neutralization reaction.

• General Equation: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

• Example: $\text{HCl} (aq) + \text{NaOH} (aq) \rightarrow \text{NaCl} (aq) + \text{H}_2\text{O} (l)$

Precipitation Reactions

Formation of Precipitates

When two aqueous solutions are mixed, an insoluble product (precipitate) may form.

• Precipitate: Solid formed from reaction in solution.

• Example: $\text{AgNO}_3 (aq) + \text{NaCl} (aq) \rightarrow \text{AgCl} (s) + \text{NaNO}_3 (aq)$

Solubility Rules Table

Solubility rules help predict whether a precipitate will form.

Compound Type	Solubility
Na+, K+, NH4+ salts	Soluble
Nitrates (NO3-), acetates (CH3COO-)	Soluble
Chlorides, bromides, iodides	Soluble (except Ag+, Pb2+, Hg22+)
Sulfates (SO42-)	Soluble (except Ba2+, Pb2+, Ca2+)
Carbonates, phosphates, sulfides	Insoluble (except Na+, K+, NH4+)
Hydroxides	Insoluble (except Na+, K+, Ba2+)

Concentration and Molarity

Definition of Molarity

Molarity (M) is a measure of concentration, defined as moles of solute per liter of solution.

• Equation: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Example: 1 mole NaCl dissolved in 1 L water = 1 M NaCl solution.

Dilution Calculations

To dilute a solution, solvent is added to decrease concentration. The number of moles of solute remains constant.

• Equation: $M_1 V_1 = M_2 V_2$

• Example: To prepare 0.5 L of 0.1 M solution from 1 M stock: $V_1 = \frac{M_2 V_2}{M_1} = \frac{0.1 \times 0.5}{1} = 0.05$ L

Stoichiometry in Solution Reactions

General Steps for Stoichiometric Calculations

Stoichiometry allows calculation of reactant and product quantities in chemical reactions.

• 1. Write balanced chemical equation.

• 2. Convert given quantities to moles.

• 3. Use mole ratios from the equation.

• 4. Convert moles to desired units (mass, volume, etc.).

• Example: How many grams of CaCO3 are needed to react with excess HCl?

• Equation: $\text{CaCO}_3 (s) + 2\text{HCl} (aq) \rightarrow \text{CaCl}_2 (aq) + \text{CO}_2 (g) + \text{H}_2\text{O} (l)$

Exchanging Ions and Net Ionic Equations

Writing Net Ionic Equations

Net ionic equations show only the species that actually participate in the reaction.

• 1. Write the balanced molecular equation.

• 2. Write the complete ionic equation (show all strong electrolytes as ions).

• 3. Cancel spectator ions to get the net ionic equation.

• Example: $\text{AgNO}_3 (aq) + \text{NaCl} (aq) \rightarrow \text{AgCl} (s) + \text{NaNO}_3 (aq)$

• Net Ionic Equation: $\text{Ag}^+ (aq) + \text{Cl}^- (aq) \rightarrow \text{AgCl} (s)$

Summary Table: Types of Electrolytes

Type	Definition	Examples
Strong Electrolyte	Completely dissociates in water	NaCl, HCl, KOH
Weak Electrolyte	Partially dissociates in water	CH3COOH, NH3
Nonelectrolyte	Does not dissociate in water	Sucrose, ethanol

Additional info:

• Some context and examples were inferred to clarify fragmented notes and ensure completeness.

• Solubility rules and electrolyte classification tables were reconstructed for clarity.