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}3\text{)}2(aq) + 2 \text{KI}(aq) \rightarrow \text{PbI}2(s) + 2 \text{KNO}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}3\text{COOH}, \text{NH}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}3)2 and KI

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

  2. Ionic Equation: \text{Pb}^{2+}(aq) + 2 \text{NO}3^{-}(aq) + 2 \text{K}^{+}(aq) + 2 \text{I}^{-}(aq) \rightarrow \text{PbI}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}2\text{H}3\text{O}2(aq) + \text{H}2\text{O} \rightarrow \text{H}3\text{O}^{+}(aq) + \text{C}2\text{H}3\text{O}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}4 , HCl, \text{H}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}2\text{H}3\text{O}2(aq) \rightleftharpoons \text{H}3\text{O}^{+}(aq) + \text{C}2\text{H}3\text{O}_2^{-}(aq)

General Ionization Equations

  • Strong Acids: \text{HX}(aq) + \text{H}2\text{O} \rightarrow \text{H}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}2\text{O} \rightleftharpoons \text{H}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}2\text{SO}3).

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

Stepwise Ionization Example for \text{H}3\text{PO}4

  1. \text{H}3\text{PO}4(aq) + \text{H}2\text{O} \rightarrow \text{H}3\text{O}^{+}(aq) + \text{H}2\text{PO}4^{-}(aq)

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

  3. \text{HPO}4^{2-}(aq) + \text{H}2\text{O} \rightleftharpoons \text{H}3\text{O}^{+}(aq) + \text{PO}4^{3-}(aq)

Acidic and Basic Anhydrides

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

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