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

Solutions and Aqueous Solutions

Definitions and Properties

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

• Solution: Homogeneous mixture of substances.

• 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.

4.1 Aqueous Solutions (aq)

Solvation and Dissolution

All substances dissolve by solvation, the process in which solvent molecules surround solute particles. Substances dissolve in water in different ways:

• Ionic compounds: Dissolve by dissociation, where water surrounds and separates the ions.

• Molecular compounds: Disperse in water, but most remain intact. Some molecular substances form ions when they dissolve.

• Example: NaCl dissociates into Na+ and Cl- ions in water.

Electrolytes

Types and Conductivity

Electrolytes are substances that produce ions in solution and thus conduct electricity. They are classified as strong, weak, or nonelectrolytes:

• Strong electrolyte: Dissociates completely in water; solution conducts electricity well.

• Weak electrolyte: Dissociates partially; solution conducts electricity poorly.

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

• Example: NaCl is a strong electrolyte; acetic acid (CH3COOH) is a weak electrolyte; sucrose is a nonelectrolyte.

Strong versus Weak Electrolytes—Equilibrium

Dissociation and Chemical Equilibrium

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

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

• Weak electrolyte equation: $\mathrm{CH_3COOH(aq) \rightleftharpoons H^+(aq) + CH_3COO^-(aq)}$

Electrolytes and Nonelectrolytes

Definitions and Examples

• Electrolyte: Substance that dissociates into ions in water (e.g., NaCl, CH3COOH).

• Nonelectrolyte: Substance that dissolves but does not form ions (e.g., glucose, sucrose).

• Example: $\mathrm{NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)}$; $\mathrm{C_6H_{12}O_6(s) \rightarrow C_6H_{12}O_6(aq)}$

4.2 Precipitation Reactions

Formation of Precipitates

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

• Precipitate: The solid formed in a precipitation reaction.

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

Solubility of Ionic Compounds

Not all ionic compounds dissolve in water. Solubility rules help predict which combinations of ions will form soluble or insoluble compounds.

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

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

Predicting Precipitate Formation

To predict whether a precipitate forms when strong electrolytes are mixed:

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.

Example: $\mathrm{Mg(NO_3)_2(aq) + 2NaOH(aq) \rightarrow Mg(OH)_2(s) + 2NaNO_3(aq)}$

Metathesis (Exchange) Reactions

Definition and Process

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

• $\mathrm{AX + BY \rightarrow AY + BX}$

• Example: $\mathrm{Mg(NO_3)_2(aq) + 2NaOH(aq) \rightarrow Mg(OH)_2(s) + 2NaNO_3(aq)}$

Completing and Balancing Metathesis Equations

1. Identify ions present in reactants.

2. Write formulas for products by combining cations and anions.

3. Check solubility rules for precipitate formation.

4. Balance the equation.

Procedure to Derive the Net Ionic Equation

Steps and Definitions

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 species involved in the reaction.

Example:

• Molecular: $\mathrm{Pb(NO_3)_2(aq) + 2KCl(aq) \rightarrow PbCl_2(s) + 2KNO_3(aq)}$

• Complete ionic: $\mathrm{Pb^{2+}(aq) + 2NO_3^-(aq) + 2K^+(aq) + 2Cl^-(aq) \rightarrow PbCl_2(s) + 2K^+(aq) + 2NO_3^-(aq)}$

• Net ionic: $\mathrm{Pb^{2+}(aq) + 2Cl^-(aq) \rightarrow PbCl_2(s)}$

4.3 Acids, Bases, and Neutralization

Acids

Acids are substances that ionize in aqueous solution to form hydrogen ions (H+). They are often called proton donors.

• Examples: Hydrochloric acid (HCl), Nitric acid (HNO3), Sulfuric acid (H2SO4), Acetic acid (CH3COOH)

Bases

Bases are substances that react with, or accept, H+ ions, increasing the concentration of hydroxide ions (OH-) in solution. Not all bases contain OH-; for example, ammonia (NH3) is a base.

• Example reaction: $\mathrm{H_2O + NH_3 \rightarrow OH^- + NH_4^+}$

Strong or Weak—Acids and Bases

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

Strong Acids	Strong Bases
Hydrochloric acid, HCl	Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH)
Hydrobromic acid, HBr	Heavy group 2A metal hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2)
Hydroiodic acid, HI
Chloric acid, HClO3
Perchloric acid, HClO4
Nitric acid, HNO3
Sulfuric acid (first proton), H2SO4

Strong or Weak Electrolyte?

To determine if a substance is a strong or weak electrolyte:

1. Is the substance ionic or molecular?

2. If ionic, it is an electrolyte. If molecular, is it an acid or base?

3. If acid: starts with H or ends in COOH. If not on the strong acid list, it is a weak acid.

4. If base: strong bases are strong electrolytes; NH3 is a weak base.

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

Neutralization Reactions

Acid-Base Reactions

Neutralization reactions occur when an acid reacts with a base, producing water and a salt (ionic compound). These reactions can be represented as molecular, complete ionic, or net ionic equations.

• Molecular: $\mathrm{HCl(aq) + NaOH(aq) \rightarrow H_2O(l) + NaCl(aq)}$

• Complete ionic: $\mathrm{H^+(aq) + Cl^-(aq) + Na^+(aq) + OH^-(aq) \rightarrow H_2O(l) + Na^+(aq) + Cl^-(aq)}$

• Net ionic: $\mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)}$

Neutralization Reactions with Gas Formation

Some metathesis reactions produce gases. For example, when a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water.

• Example: $\mathrm{CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)}$

• Example: $\mathrm{NaHCO_3(aq) + HBr(aq) \rightarrow NaBr(aq) + CO_2(g) + H_2O(l)}$

Reactions of sulfides with acids produce hydrogen sulfide gas (H2S), which is poisonous and has a characteristic odor.

• Example: $\mathrm{Na_2S(aq) + H_2SO_4(aq) \rightarrow Na_2SO_4(aq) + H_2S(g)}$

Displacement Reactions

Oxidation of Metals by Acids and Salts

In displacement reactions, ions oxidize an element, often resulting in the release of a gas. For example, H+ ions oxidize Mg(s), producing H2(g).

• Example: $\mathrm{Mg(s) + 2H^+(aq) \rightarrow Mg^{2+}(aq) + H_2(g)}$

4.5 Concentration and Molarity

Measuring Solution Concentration

The concentration of a solution is the amount of solute dissolved in a given quantity of solvent. Molarity (M) is a common unit of concentration, defined as:

• $\mathrm{Molarity\ (M) = \frac{\text{moles of solute}}{\text{volume of solution in liters}}}$

• Molarity can be used as a conversion factor between moles and liters.

Mixing a Solution

To prepare a solution of known molarity:

1. Weigh out a known mass (number of moles) of solute.

2. Add solute to a volumetric flask.

3. Add solvent to the calibration line on the flask neck.

Dilution

A solution can be diluted by adding only solvent, lowering the concentration but not changing the number of moles of solute.

• The relationship between concentrated and dilute solutions is given by:

$\mathrm{M_c \times V_c = M_d \times V_d}$

• Where $M_c$ and $V_c$ are the molarity and volume of the concentrated solution, and $M_d$ and $V_d$ are those of the dilute solution.

4.6 Solution Stoichiometry and Chemical Analysis

Stoichiometric Calculations

Stoichiometry in solution involves using molarity and volume to relate quantities of reactants and products.

• Convert mass of substance to moles using molar mass.

• Use coefficients from balanced equations to relate moles of different substances.

• Use molarity and volume to find moles in solution.

Titration

Titration is an analytical technique used to determine the concentration of a solute in solution. A solution of known concentration (standard solution) is added to a solution of unknown concentration until the reaction reaches the equivalence point, often indicated by a color change (end point).

• Standard solution: Solution of known concentration.

• Equivalence point: Point at which stoichiometric amounts of reactants have reacted.

• End point: Observable change (e.g., color) indicating completion.

Summary Tables

Solubility Guidelines for Common Ionic Compounds in Water

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

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

Common Strong Acids and Bases

Strong Acids	Strong Bases
HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4 (first proton)	LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Electrolytic Behavior of Common Soluble Ionic and Molecular Compounds

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

Additional info: These notes expand on the provided slides and text, offering definitions, examples, and equations for key concepts in Chapter 4: Reactions in Aqueous Solution, suitable for General Chemistry students.