CH 4.1

Chapter 4: Molecular View of Reactions in Aqueous Solutions Part I

Chapter Objectives and Learning Goals

• Describe solutions qualitatively and quantitatively.

• Distinguish electrolytes from non-electrolytes.

• Write balanced molecular, ionic, and net ionic equations.

• Identify acids and bases, including their names and formulas.

• Use metathesis reactions to plan chemical syntheses.

• Define and use molarity in calculations.

• Understand titrations and chemical analysis.

Importance of Water

• Common Compound: Abundant on Earth.

• Solvent Properties: Dissolves many substances, vital for life.

• Biological Relevance: Approximately 60% of the human body.

• Role as a Solvent: Dissolves ionic compounds and facilitates acid-base reactions.

Reactions in Solution

• Physical Contact: Reactants must physically contact (gas or liquid phases).

• Definition of a Solution: A homogeneous mixture where components mix freely.

Definitions

Solvent

• The medium that dissolves solutes, present in the largest amount.

  • Aqueous Solution: Water is the solvent.

Solute

• Substance dissolved in a solvent.

  • Examples: Gas ($\text{CO}_2$ in soda), Liquid (Ethylene glycol), Solid (Sugar).

Concentration of Solutions

• Characterization: Solute-to-solvent ratio, defined in grams per gram.

  • Percent Concentration: $\text{g solute} \times \frac{100}{\text{g solution}}$

Relative Concentration

• Dilute Solution: Small solute-to-solvent ratio.

• Concentrated Solution: Large solute-to-solvent ratio.

• Terms are relative.

Concentration and Solubility

• Solubility: Amount of solute (g) to create a saturated solution at a specific temperature: $\text{Solubility} = \frac{\text{g solute}}{100 \text{ g solvent}}$

• Temperature Dependence: Solubility varies with temperature.

  • Saturated Solution: Maximum solute dissolved at a given temperature.

  • Unsaturated Solution: Less than maximum capacity; more can dissolve.

Common Substance Solubility Example

• Sodium chloride (NaCl): 35.7 g/100g water at 0 °C; 39.1 g/100g water at 100 °C.

Supersaturated Solutions

• Definition: Contains more solute than saturation at a given temperature.

• Formation and Stability: Formed by cooling saturated solutions; unstable, crystallize upon adding a seed crystal (precipitate formation).

Precipitates

Definition of a Precipitate

• Solid formed in solutions when a product has low solubility.

• Insoluble Product: Separates from the solution.

Precipitation Reactions

• Example Reaction: $\text{Pb(NO}<em>3\text{)}</em>2(aq) + 2 \text{KI}(aq) \rightarrow \text{PbI}<em>2(s) + 2 \text{KNO}</em>3(aq)$

Electrolytes in Aqueous Solution

• Definition of Electrolytes: Ionic compounds that conduct electricity when dissolved in water.

Ionic Compounds in Water

• Dissociation Process: Water molecules surround and remove ions from the lattice (e.g., NaCl into $\text{Na}^{+}$ and $\text{Cl}^{-}$). Polyatomic ions stay intact.

Molecular Compounds in Water

• Dissolution Behavior: Molecules dissolve but do not dissociate into ions.

Electrical Conductivity

Types of Electrolytes

• Strong Electrolyte: Completely dissociates in water (e.g., NaCl, $\text{HClO}_4$).

• Weak Electrolyte: Partially dissociates (small percentage ionize) (e.g., $\text{CH}<em>3\text{COOH}$, $\text{NH}</em>3$).

• Non-Electrolyte: Does not conduct electricity (intact molecular structure) (e.g., sugar, alcohols).

Dissociation Reactions

• Process Description: Ionic compounds dissolve to form hydrated ions, indicated by (aq).

• Example Dissociation Reaction Formula: $\text{KBr}(s) \rightarrow \text{K}^{+}(aq) + \text{Br}^{-}(aq)$

Equations of Ionic Reactions

Types of Equations

1. Molecular Equation: Substances listed with complete formulas.

2. Ionic Equation: All soluble substances broken into ions.

3. Net Ionic Equation: Only species undergoing chemical change.

Example of Reaction Equations between Pb($\text{NO}<em>3$)$</em>2$ and KI

1. Molecular Equation: $\text{Pb(NO}<em>3\text{)}</em>2(aq) + 2 \text{KI}(aq) \rightarrow \text{PbI}<em>2(s) + 2 \text{KNO}</em>3(aq)$

2. Ionic Equation: $\text{Pb}^{2+}(aq) + 2 \text{NO}<em>3^{-}(aq) + 2 \text{K}^{+}(aq) + 2 \text{I}^{-}(aq) \rightarrow \text{PbI}</em>2(s) + 2 \text{K}^{+}(aq) + 2 \text{NO}_3^{-}(aq)$

3. Net Ionic Equation: $\text{Pb}^{2+}(aq) + 2 \text{I}^{-}(aq) \rightarrow \text{PbI}_2(s)$

Spectator Ions

• Definition: Ions that do not participate in the chemical reaction.

Criteria for Balancing Ionic and Net Ionic Equations

1. Material Balance: Same number of each atom on both sides.

2. Electrical Balance: Net electrical charge equal on both sides.

Acids and Bases

• Common Laboratory Reagents: Acids (tart/sour taste like vinegar) and bases (bitter taste, slippery feel like ammonia). Safety Note: Never taste, feel, or smell chemicals.

Definitions of Acids and Bases

Arrhenius Acid

• Definition: Produces hydronium ions ($\text{H}_3\text{O}^{+}$) in water.

• Example: $\text{HC}<em>2\text{H}</em>3\text{O}<em>2(aq) + \text{H}</em>2\text{O} \rightarrow \text{H}<em>3\text{O}^{+}(aq) + \text{C}</em>2\text{H}<em>3\text{O}</em>2^{-}(aq)$

Arrhenius Base

• Definition: Produces hydroxide ions ($\text{OH}^{-}$).

• Example: $\text{NaOH}(s) \rightarrow \text{Na}^{+}(aq) + \text{OH}^{-}(aq)$

Strong Acids

• Characteristics: Fully dissociate in water (e.g., $\text{HClO}<em>4$ , HCl, $\text{H}</em>2\text{SO}_4$).

Strong Bases

• Characteristics: Completely dissociate in water (e.g., Group 1A & 2A metal hydroxides like LiOH, NaOH).

Weak Acids and Bases

• Weak Acids: < 100% ionization (e.g., $\text{CH}_3\text{COOH}$).

• Weak Bases: Molecular bases that don't dissociate efficiently (e.g., $\text{NH}_3$).

Dynamic Equilibrium

• Definition: Opposing reactions at the same rate, constant concentrations.

• Example: $\text{HC}<em>2\text{H}</em>3\text{O}<em>2(aq) \rightleftharpoons \text{H}</em>3\text{O}^{+}(aq) + \text{C}<em>2\text{H}</em>3\text{O}_2^{-}(aq)$

General Ionization Equations

• Strong Acids: $\text{HX}(aq) + \text{H}<em>2\text{O} \rightarrow \text{H}</em>3\text{O}^+(aq) + X^-(aq)$

• Strong Bases: $\text{M(OH)}_n(s) \rightarrow \text{M}^{n+}(aq) + n \text{OH}^-(aq)$

• Weak Acids: $\text{HA}(aq) + \text{H}<em>2\text{O} \rightleftharpoons \text{H}</em>3\text{O}^+(aq) + \text{A}^-(aq)$

• Weak Bases: $\text{B}(aq) + \text{H}_2\text{O} \rightleftharpoons \text{HB}^+(aq) + \text{OH}^-(aq)$

Summary of Electrolytes

• Strong Electrolytes: Completely ionize, predominantly products.

• Weak Electrolytes: Partially ionize, significant reverse reactions (more reactants).

Polyprotic Acids

Characteristics

• Monoprotic Acids: Furnish one $\text{H}^{+}$ (e.g., $\text{HNO}_3$).

• Diprotic Acids: Furnish two $\text{H}^{+}$ (e.g., $\text{H}<em>2\text{SO}</em>3$).

• Triprotic Acids: Furnish three $\text{H}^{+}$ (e.g., $\text{H}<em>3\text{PO}</em>4$).

Stepwise Ionization Example for $\text{H}<em>3\text{PO}</em>4$

1. $\text{H}<em>3\text{PO}</em>4(aq) + \text{H}<em>2\text{O} \rightarrow \text{H}</em>3\text{O}^{+}(aq) + \text{H}<em>2\text{PO}</em>4^{-}(aq)$

2. $\text{H}<em>2\text{PO}</em>4^{-}(aq) + \text{H}<em>2\text{O} \rightleftharpoons \text{H}</em>3\text{O}^{+}(aq) + \text{HPO}_4^{2-}(aq)$

3. $\text{HPO}<em>4^{2-}(aq) + \text{H}</em>2\text{O} \rightleftharpoons \text{H}<em>3\text{O}^{+}(aq) + \text{PO}</em>4^{3-}(aq)$

Acidic and Basic Anhydrides

• Acidic Anhydrides: Nonmetal oxides reacting with water to form acids (e.g., $\text{SO}<em>3 + \text{H}</em>2\text{O} \rightarrow \text{H}<em>2\text{SO}</em>4$).

• Basic Anhydrides: Metal oxides reacting with water to form bases (e.g., $\text{CaO} + \text{H}<em>2\text{O} \rightarrow \text{Ca(OH)}</em>2$).

Nomenclature of Acids and Bases

Naming Acids

• Binary Acids: Hydrogen and a nonmetal (e.g., HCl [g] $\rightarrow$ hydrochloric acid [aq]).

• Oxoacids: Hydrogen, oxygen, and another non-metal, named based on parent oxoanion.

Naming Bases

• Metal Hydroxides: Named by the metal (e.g., $\text{Ca(OH)}_2 \rightarrow \text{calcium hydroxide}$).

• Molecular Bases: Named like organic molecules (e.g., $\text{NH}_3$ is ammonia).