Chapter 4: Reactions in Aqueous Solution – General Chemistry Study Notes

Reactions in Aqueous Solution

Solutions

Solutions are homogeneous mixtures composed of two or more pure substances. The solvent is the component present in the greatest amount, while all other components are called solutes. When water acts as the solvent, the mixture is termed an aqueous solution.

• Homogeneous mixture: Uniform composition throughout.

• Solvent: Substance in greatest abundance.

• Solute: Substance(s) dissolved in the solvent.

• Aqueous solution: Solution where water is the solvent.

• Example: Salt water, where water is the solvent and salt is the solute.

Electrolytes

Electrolytes are substances that produce ions when dissolved in water, enabling the solution to conduct electricity. Their behavior is classified as strong, weak, or nonelectrolyte.

• Strong electrolyte: Dissociates completely in water; strong conductor of electricity.

• Weak electrolyte: Dissociates partially; weak conductor.

• Nonelectrolyte: Does not dissociate; does not conduct electricity.

• Example: NaCl is a strong electrolyte; sucrose is a nonelectrolyte.

4.1 Aqueous Solutions (aq)

Substances dissolve in water by solvation, where solvent molecules surround solute particles. Ionic compounds dissolve by dissociation, forming separated ions, while molecular compounds may disperse but often remain intact. Some molecular substances can form ions upon dissolution.

• Solvation: Surrounding of solute by solvent molecules.

• Dissociation: Separation of ions in ionic compounds.

• Example: NaCl(s) → Na+(aq) + Cl-(aq)

Electrolytes and Nonelectrolytes

An electrolyte dissociates into ions in water (e.g., NaCl, CH3COOH). A nonelectrolyte dissolves but does not form ions (e.g., glucose, sucrose).

• Electrolyte examples: NaCl(s) → Na+(aq) + Cl-(aq)

• Nonelectrolyte examples: C6H12O6(s) → C6H12O6(aq)

Strong versus Weak Electrolytes—Equilibrium

Strong electrolytes dissociate completely, indicated by a single arrow in equations. Weak electrolytes only partially dissociate, represented by a double arrow, indicating chemical equilibrium.

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

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

Important Points

• Water-soluble ionic compounds are strong electrolytes.

• Ionic compounds usually consist of a metal and a non-metal.

• Ammonium salts (e.g., NH4Br) are ionic, even without a metal.

• Acids and bases ionize in water.

4.2 Precipitation Reactions

Precipitation reactions occur when two solutions of soluble salts are mixed and an insoluble salt (precipitate) forms.

• Precipitate: The solid formed in the reaction.

• Example: Mixing AgNO3(aq) and NaCl(aq) forms AgCl(s).

Solubility of Ionic Compounds

Not all ionic compounds dissolve in water. Solubility less than 0.01 mol/L is considered insoluble. Solubility rules help predict which combinations will dissolve.

Soluble Ionic Compounds	Important Exceptions
NO3-	None
CH3COO-	None
Cl-	Ag+, Hg22+, Pb2+
Br-	Ag+, Hg22+, Pb2+
I-	Ag+, Hg22+, Pb2+
SO42-	Sr2+, Ba2+, Hg22+, Pb2+

Insoluble Ionic Compounds	Important Exceptions
S2-	NH4+, alkali metals, Ca2+, Sr2+, Ba2+
CO32-	NH4+, alkali metals
PO43-	NH4+, alkali metals
OH-	NH4+, alkali metals, Ca2+, Sr2+, Ba2+

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

• Ionic compounds with nitrate (NO3-) are soluble.

How to Predict Whether a Precipitate Forms

1. Note the ions present in the reactants.

2. Consider possible cation-anion combinations.

3. Use solubility rules to determine if any combination is insoluble.

Metathesis (Exchange) Reactions

Metathesis reactions involve the exchange of ions between reactants. The general form is:

• $AX + BY \rightarrow AY + BX$

• Example: $\text{Mg(NO}_3)_2(aq) + 2\text{NaOH}(aq) \rightarrow \text{Mg(OH)}_2(s) + 2\text{NaNO}_3(aq)$

Completing and Balancing Metathesis Equations

1. Identify ions in reactants.

2. Write product formulas using cation from one and anion from the other.

3. Check solubility rules for precipitate formation.

4. Balance the equation.

Procedure to Derive the Net Ionic Equation

1. Molecular equation: Balanced equation showing reactants and products.

2. Complete ionic equation: Shows all strong electrolytes as ions.

3. Net ionic equation: Cancels spectator ions, showing only those involved in forming the product.

Molecular, Complete Ionic, and Net Ionic Equations

• Molecular equation: $\text{Pb(NO}_3)_2(aq) + 2\text{KCl}(aq) \rightarrow \text{PbCl}_2(s) + 2\text{KNO}_3(aq)$

• Complete ionic equation: $\text{Pb}^{2+}(aq) + 2\text{NO}_3^-(aq) + 2\text{K}^+(aq) + 2\text{Cl}^-(aq) \rightarrow \text{PbCl}_2(s) + 2\text{K}^+(aq) + 2\text{NO}_3^-(aq)$

• Net ionic equation: $\text{Pb}^{2+}(aq) + 2\text{Cl}^-(aq) \rightarrow \text{PbCl}_2(s)$

4.3 Acids, Bases, and Neutralization

Acids ionize in water to form H+ ions and are called proton donors. Bases accept H+ ions or increase OH- concentration. Not all bases contain OH- (e.g., NH3).

• Strong acids: Completely dissociate (e.g., HCl, HNO3).

• Weak acids: Partially dissociate (e.g., CH3COOH).

• Strong bases: Group 1A and 2A metal hydroxides (e.g., NaOH, Ca(OH)2).

• Weak bases: Partially react (e.g., NH3).

Strong Acids	Strong Bases
HCl, HBr, HI, HClO4, HNO3, H2SO4	LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Strong or Weak Electrolyte?

• If ionic, it is an electrolyte.

• If molecular, check if it is an acid or base.

• Strong acids and bases are strong electrolytes; weak acids and bases are weak electrolytes.

	Strong Electrolyte	Weak Electrolyte	Nonelectrolyte
Ionic	All	None	None
Molecular	Strong acids	Weak acids, weak bases	All other compounds

Neutralization Reactions

Neutralization occurs when an acid reacts with a base, producing water and a salt. Equations can be written in molecular, complete ionic, or net ionic forms.

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

• Net ionic: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$

Neutralization Reactions with Gas Formation

• Carbonates/bicarbonates with acids produce salt, CO2, and water.

• Sulfides with acids produce H2S gas (poisonous, rotten egg odor).

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

4.4 Oxidation-Reduction Reactions

Redox reactions involve the transfer of electrons. Oxidation is the loss of electrons, while reduction is the gain of electrons. Both processes occur simultaneously.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Redox reaction: Reaction involving electron transfer.

Oxidation Numbers

Oxidation numbers are assigned to elements in compounds to track electron transfer. They do not represent actual charges but are a bookkeeping tool.

• Atoms in elemental form: oxidation number = 0.

• Monatomic ions: oxidation number = ion charge.

• Nonmetals: usually negative, but can be positive in certain compounds.

• Sum of oxidation numbers in a neutral compound = 0; in a polyatomic ion = ion charge.

Rules to Assign Oxidation Numbers

1. Elemental form: 0 (e.g., Na(s), O2(g)).

2. Monatomic ion: equals charge (e.g., Na+ = +1).

3. Oxygen: usually -2, except in peroxides (-1).

4. Hydrogen: +1 (with nonmetals), -1 (with metals).

5. Fluorine: always -1; other halogens usually -1 unless with oxygen.

6. Sum in neutral compound = 0; sum in polyatomic ion = ion charge.

Sample Exercises

• Relating numbers of anions and cations to chemical formulas.

• Writing net ionic equations for precipitation reactions.

• Identifying strong, weak, and nonelectrolytes by conductivity.

• Solution stoichiometry and chemical analysis using dilutions and titrations.

Additional info: These notes cover the essential concepts and procedures for understanding reactions in aqueous solution, including solution composition, electrolytes, precipitation, acid-base and redox reactions, and practical applications such as antacids.