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 is an aqueous solution of NaCl in H2O.

4.1 Aqueous Solutions (aq)

Solvation and Dissolution Mechanisms

All substances dissolve by solvation, where solute particles are surrounded by solvent molecules. The manner in which substances dissolve in water depends on their chemical nature:

• Ionic compounds: Dissolve by dissociation, separating into ions surrounded by water molecules.

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

• Example: NaCl(s) → Na+(aq) + Cl-(aq); CH3COOH(aq) → CH3COO-(aq) + H+(aq)

Electrolytes

Types and Conductivity

Electrolytes are substances that produce ions in solution, enabling electrical conductivity. Their strength depends on the degree of dissociation:

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

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

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

• Example: NaCl is a strong electrolyte; CH3COOH is a weak electrolyte; glucose is a nonelectrolyte.

Strong versus Weak Electrolytes—Equilibrium

Chemical Equilibrium in Dissolution

The distinction between strong and weak electrolytes is reflected in chemical equations:

• Strong electrolyte: Complete dissociation, indicated by a single arrow ($\rightarrow$).

• Weak electrolyte: Partial dissociation, indicated by a double arrow ($\rightleftharpoons$), representing chemical equilibrium.

• Example Equations:

$\mathrm{HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)}$ $\mathrm{CH_3COOH(aq) \rightleftharpoons H^+(aq) + CH_3COO^-(aq)}$

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: C6H12O6(s) → C6H12O6(aq)

4.2 Precipitation Reactions

Formation of Insoluble Salts

Precipitation reactions occur when two solutions containing soluble salts are mixed and an insoluble salt, called a precipitate, forms.

• Precipitate: The solid formed in a precipitation reaction.

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

Solubility of Ionic Compounds

Solubility Rules and Guidelines

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

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+

How to Predict Whether a Precipitate Forms

Steps for Prediction

• 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 General Form

Metathesis reactions involve the exchange of ions between reactant 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

Stepwise Procedure

• 1) Identify ions present in reactants.

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

• 3) Apply solubility rules to identify precipitates.

• 4) Balance the equation.

Procedure to Derive the Net Ionic Equation

Types of Equations

• Molecular equation: Shows reactants and products as compounds.

• Complete ionic equation: Shows all strong electrolytes as dissociated ions.

• Net ionic equation: Shows only the species that actually participate in the reaction; spectator ions are omitted.

• 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: Substances that ionize in aqueous solution to form H+ ions (proton donors).

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

Bases

• Bases: Substances that react with or accept H+ ions, increasing the concentration of OH- ions 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—Acid and Bases

Classification and Examples

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

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

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

• Weak bases: Partially react to produce hydroxide ions (e.g., NH3).

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?

Identification Steps

• 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.

• Equations can be written as molecular, complete ionic, or net ionic forms.

• Example:

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

Special Cases

• When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water.

• When sulfides react with acids, hydrogen sulfide gas (H2S) is produced, which is poisonous and has a rotten egg odor.

• Examples:

$\mathrm{CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)}$ $\mathrm{Na_2S(aq) + H_2SO_4(aq) \rightarrow Na_2SO_4(aq) + H_2S(g)}$

Displacement Reactions

Oxidation of Metals by Acids and Salts

• In displacement reactions, ions oxidize an element, often producing hydrogen gas.

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

4.5 Concentration and Molarity

Measuring Solution Concentration

• Concentration: Amount of solute dissolved in a given quantity of solvent.

• Molarity (M): Defined as moles of solute per liter of solution.

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

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

Mixing a Solution

Preparation of Standard Solutions

• Weigh out a known mass of solute to determine moles.

• Add solute to a volumetric flask and add solvent up to the calibration mark.

Dilution

Lowering Concentration

• A solution can be diluted by adding only solvent; the number of moles of solute remains unchanged.

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

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

• Where $M_c$ and $M_d$ are the molarities, and $V_c$ and $V_d$ are the volumes of the concentrated and dilute solutions, respectively.

4.6 Solution Stoichiometry and Chemical Analysis

Stoichiometric Calculations

• Use molarity and volume to determine moles of solute.

• Apply coefficients from balanced equations to relate moles of reactants and products.

• Convert between mass, moles, and volume as needed.

Titration

Analytical Technique for Concentration Determination

• Titration: Technique to determine the concentration of a solute in solution using a standard solution.

• The reaction is complete at the equivalence point, often indicated by a color change (end point).

• Standard solution: Solution of known concentration used in titration.

Summary Table: Electrolytic Behavior of Compounds

Type	Strong Electrolyte	Weak Electrolyte	Nonelectrolyte
Ionic	All	None	None
Molecular	Strong acids/bases	Weak acids/bases	All other compounds

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