Chapter 7: Reactions in Aqueous Solution

7.1 Introduction

Solution: Any homogeneous mixture that is physically and chemically the same throughout the whole system.
- Two main components:
  - Solvent: Component present in large amounts (typically water in aqueous solutions).
  - Solute: The substance that is dissolved in the solvent; usually present in relatively small amounts (in moles).
- Key characteristics: There can be more than one solute in a solution, but there is never more than one solvent.

7.2 Solubility

Aqueous solubility: Solutes dissolve in water to form a homogeneous mixture;
- Water: Known as the universal solvent.
- High solubility: Ionic solids and some polar compounds, such as strong acids and strong bases, tend to have the highest solubility. They form ions when dissolved in water.
- Terminology:
  - Soluble/insoluble decisions regarding ionic solids.
  - Miscible/immiscible distinctions for liquid solutes.

7.2.1 Electrolyte Solutions

Definition: An electrolyte solution is any aqueous solution that conducts electricity.
- Examples of strong electrolytes:
  - Ionic compounds such as sodium chloride (NaCl).
  - Strong electrolytes dissociate completely into their constituent ions.
    - Dissociation of NaCl:
      ext{NaCl(aq) }
      ightarrow ext{ Na}^+(aq) + ext{ Cl}^-(aq)
  - Water's polarity aids in ion-dipole interactions that facilitate this dissociation:
    - The − end of the water molecule is attracted to Na+; the + end is attracted to Cl−.
  - Other examples of strong electrolytes: Potassium hydroxide (KOH) which dissociates as follows:
    ext{NaOH(aq) }
    ightarrow ext{ Na}^+(aq) + ext{ OH}^-(aq)
- Weak electrolytes: Solutes that dissolve in water producing fewer ions, leading to lower conductivity. Examples include:
  - Weak acids like acetic acid (HC₂H₃O₂) and weak bases like ammonia (NH₃).
    - Example reaction for acetic acid:
      ext{HC}2 ext{H}3 ext{O}2(aq) ightleftharpoons ext{CH}3 ext{COO}^-(aq) + ext{H}^+(aq)
- Non-electrolytes: Substances that dissolve in water to form neutral molecules and have virtually no effect on electrical conductivity.
  - Examples: Ethanol (C₂H₆O), ethylene glycol (C₂H₆O₂), glucose (C₆H₁₂O₆), and sucrose (C₁₂H₂₂O₁₁).

7.2.2 Solubility Rules

Purpose: Determine which ionic solids dissolve in water.
- Rule #1: Cation Rule: If M (cation of ionic compound MX) is Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺, or NH₄⁺, MX is soluble in water.
  - Example: NaCl and Na₂CO₃ (both contain Na⁺) are soluble in water.
- Rule #2: Anion Groups Rule: If X (anion of ionic compound MX) is NO₃⁻ (nitrate), ClO₄⁻ (perchlorate), or CH₃CO₂⁻ (acetate), MX is soluble in water.
  - Examples: NaNO₃ and Ba(NO₃)₂ are soluble as they contain NO₃⁻.
- Rule #3: Use a solubility table for ions not covered in Rules 1 and 2.
  - Soluble Compounds:
    - Group IA cations and NH₄⁺ are soluble without exceptions.
    - Nitrate (NO₃⁻), perchlorate (ClO₄⁻), acetate (CH₃CO₂⁻) are all soluble without exceptions.
    - Halides (F⁻, Cl⁻, Br⁻, I⁻) are soluble except when paired with Ag⁺, Hg₂²⁺, Pb²⁺.
  - Insoluble Compounds:
    - Carbonates (CO₃²⁻) are insoluble except with Group IA cations and NH₄⁺.
    - Hydroxides (OH⁻) are insoluble except with Group IA cations and some earth alkali metals: Ca²⁺, Ba²⁺, Sr²⁺.

7.3 Precipitation Reactions

Definition: Reactions occur when two soluble aqueous ionic solutions react to form a precipitate (an insoluble ionic solid in water).
- Example of precipitation: ext{AgNO}3(aq) + ext{NaCl}(aq) ightarrow ext{AgCl}(s) + ext{NaNO}3(aq)
  - NaCl, AgNO₃, and NaNO₃ are completely soluble in water; AgCl is insoluble, forming a white precipitate.
Double-Displacement Reaction: Type of chemical reaction where cations and anions of two reactants switch places, forming two new compounds (products).
- Example displacements: AX + MY ightarrow MX + AY
  - Characteristics of double-replacement reactions include precipitation reactions and acid-base neutralizations.
- Three representations of reactions:
  1. Molecular Equation: Shows all reactants and products in neutral form. Example:
    ext{NiCl}2(aq) + ext{Na}2 ext{S}(aq)
    ightarrow ext{NiS}(s) + 2 ext{NaCl}(aq)
  2. Complete Ionic Equation: Shows strong electrolytes in dissociated form.
    ext{Ni}^{2+}(aq) + 2 ext{Cl}^-(aq) + 2 ext{Na}^+(aq) + ext{S}^{2-}(aq)
    ightarrow ext{NiS}(s) + 2 ext{Na}^+(aq) + 2 ext{Cl}^-(aq)
  3. Net Ionic Equation: Shows only the species that partake in the reaction, canceling common ions (spectator ions).
    - Example:
      ext{Ni}^{2+}(aq) + ext{S}^{2-}(aq)
      ightarrow ext{NiS}(s)

7.4 Acid-Base Neutralization Reactions

Arrhenius Acid: Dissolves in water to produce hydrogen ions (H⁺). Example: ext{HA}(aq) ightarrow ext{H}^+(aq) + ext{A}^−(aq)
- Example: Hydrochloric acid (HCl):
  ext{HCl}(aq)
  ightarrow ext{H}^+(aq) + ext{Cl}^−(aq)
Arrhenius Base: Dissolves in water to produce hydroxide ions (OH⁻). Example: ext{BOH}(aq) ightarrow ext{B}^+(aq) + ext{OH}^−(aq)
- Example: Sodium hydroxide (NaOH):
  ext{NaOH}(aq)
  ightarrow ext{Na}^+(aq) + ext{OH}^−(aq)
Neutralization Reactions: When an acid reacts with a base to yield a salt and water.
- General Reaction:
  ext{HA}(aq) + ext{BOH}(aq)
  ightarrow ext{AB}(aq) + ext{H}_2 ext{O}(l)
- Example:
  ext{HCl}(aq) + ext{NaOH}(aq)
  ightarrow ext{NaCl}(aq) + ext{H}_2 ext{O}(l)
Types of Acid-Base Reactions:
- Strong Acid and Strong Base:
  - Net Ionic Equation: Identifies the primary species.
    1. Molecular Equation: ext{HCl}(aq) + ext{NaOH}(aq) ightarrow ext{NaCl}(aq) + ext{H}_2 ext{O}(l)
      1. Complete Ionic Equation:
        ext{H}^+(aq) + ext{Cl}^−(aq) + ext{Na}^+(aq) + ext{OH}^−(aq)
        ightarrow ext{Na}^+(aq) + ext{Cl}^−(aq) + ext{H}_2 ext{O}(l)
      2. Net Ionic Equation:
        ext{H}^+(aq) + ext{OH}^−(aq)
        ightarrow ext{H}_2 ext{O}(l)
- Weak Acid and Strong Base:
  - Example: Neutralization between HF (weak acid) and KOH (strong base).
    - Molecular Equation:
      ext{HF}(aq) + ext{KOH}(aq)
      ightarrow ext{KF}(aq) + ext{H}_2 ext{O}(l)
    - Complete Ionic Equation:
      ext{HF}(aq) + ext{K}^+(aq) + ext{OH}^−(aq)
      ightarrow ext{K}^+(aq) + ext{F}^−(aq) + ext{H}_2 ext{O}(l)
    - Net Ionic Equation:
      ext{HF}(aq) + ext{OH}^−(aq)
      ightarrow ext{F}^−(aq) + ext{H}_2 ext{O}(l)

7.5 Redox Reactions

Definition: Electron transfer reactions (oxidation–reduction reactions) where electrons move from one species to another.
- Oxidation: Substance loses one or more electrons (OIL: Oxidation Is Loss).
- Reduction: Substance gains one or more electrons (RIG: Reduction Is Gain).
Representing Redox Reactions: For example, when magnesium burns:
ext{2Mg}(s) + ext{O}_2(g)
ightarrow ext{2MgO}(s)
General Redox Representation:
- For a chemical reaction between copper and silver ions in solution: ext{Cu}(s) + 2 ext{Ag}^+(aq) ightarrow 2 ext{Ag}(s) + ext{Cu}^{2+}(aq)
  - Each Cu atom loses 2e− (oxidation) while Ag⁺ ions gain 1e− each (reduction).
Oxidation States: Indicate the number of electrons gained or lost when atoms bond with another element.
- Guidelines for determining oxidation states include:
  - Any pure element has an oxidation state of 0.
  - Cations of Group 1 (IA +1), Group 2 (IIA +2), and Group 3 (IIIA +3).
  - Group 16(VIA) anions have oxidation states of -2; Group 17(VIIA) anions typically -1 with more metallic elements.
  - Transition metals can have multiple oxidation states.
Determining Oxidation States:
- Use algebra to assign oxidation states, and remember that in balanced equations, oxidation states remain consistent regardless of coefficients.
- Example: The oxidation state of Cr in potassium dichromate (K₂Cr₂O₇) can be determined through systematic assigning starting with K and O to find Cr's contribution as +6.

7.6 Solution Concentration

Definition: The concentration of a solution reflects the amount of solute in a given amount of solvent or solution.
- Molarity (M): Defined as the amount of solute in moles per liter of solution. Formula:
  ext{Molarity} = rac{ ext{amount of solute (moles)}}{ ext{volume of solution (liters)}}
Example Calculation: If a 355 mL soft drink sample contains 0.133 mol of sucrose, the molarity can be calculated as follows:
M = rac{0.133 ext{ mol}}{0.355 ext{ L}} = 0.375 ext{ M}
Preparing a Known Concentration Solution: Combine known solute and solvent in a volumetric flask and mix until the desired volume is reached.
Example Problem: Finding grams of K₂Cr₂O₇ needed for a 2.16 M concentration:
1. Use the formula:
  ext{moles of solute} = ext{M} imes V
2. Plug values:
  ext{moles} = 2.16 ext{ mol/L} imes 0.2500 ext{ L} = 0.540 ext{ mol}
3. Convert moles to grams using molar mass:
  0.540 ext{ mol} imes 294.2 ext{ g/mol} = 159 ext{ g}
Dilution: The process of decreasing the concentration of a solution by adding solvent.
- Dilution formula: M1 V1 = M2 V2
  - Where M is molarity and V is volume.
Example Dilution Problem: From a 5.00 M solution diluted to 1.80 L:
1. Calculate V2 using M1:
  V2 = rac{M1 V1}{M2}
  where M1 = 5.00, V1 = 0.850 L, M2 = unknown.

7.7 Solution Stoichiometry

Use of Molarity in Stoichiometric Calculations: Concentrations allow calculation of moles for reactions; e.g., if 0.225 M MgCl₂ needs to react with 0.250 M AgNO₃ to precipitate AgCl:
1. Identify limiting reactants and determine volume required via stoichiometric relationships:
- Given reaction:
  2 ext{AgNO}3 (aq) + ext{MgCl}2(aq)
  ightarrow ext{Mg(NO}3)2(aq) + 2 ext{AgCl}(s)
- Calculate the needed volume based on concentration ratios for complete precipitation.