Chapter 4: Aqueous Reactions - Solubility and Reaction Types

Course Overview and Resources

  • This week (following the exam-free previous week) covers Chapter 4: Aqueous Reactions, and begins Chapter 5: Thermochemistry on Friday.

  • Available resources for mastering Chapter 4 material include office hours, supplemental instruction (SI), Undergraduate Teaching Fellows (UTF), and tutoring.

  • Recap from Friday: Introduction to aqueous solutions, electrolytes vs. non-electrolytes, and electrolytes include salts, acids, and bases.

Solubility

  • Definition: Solubility refers to how much of a solute can dissolve in a given amount of solvent.

    • Example: Adding sugar to water initially dissolves, but eventually reaches a saturation point where the solution is saturated, and no more solute can dissolve.

  • Reporting Solubility: The maximum amount of solute that can dissolve can be reported in units such as "x grams of sucrose per 100 ml of solution."

Degrees of Solubility

  1. Soluble: A significant amount can dissolve.

    • Examples: Glucose, sucrose, and certain ionic compounds like sodium chloride (NaCl) and magnesium sulfate (MgSO4), certain bases like sodium hydroxide (NaOH), and certain acids like acetic acid (CH3COOH).

  2. Insoluble: Does not dissolve very well in water.

    • Examples: Calcium carbonate (CaCO_3) (responsible for hard water deposits), silver chloride (AgCl), hexane, and iodine.

    • Note: Some ionic compounds are soluble, others are not.

  3. Miscible: Can dissolve in any proportions whatsoever.

    • Unlike sucrose, which reaches a saturation point, miscible liquids can be mixed indefinitely.

    • Examples: Water with ethanol, glycerol, or formaldehyde.

Solubility of Salts

  • Definition of Salt: An ionic compound containing a metal cation (or polyatomic ion like ammonium, NH_4^+) and a partner anion (non-metal or polyatomic ion).

  • General Rule: Most salts are actually soluble.

  • Distinction: The difference between soluble and insoluble is based on the quantity that can dissolve (a "significant amount").

  • Instructor's Policy: Students will not be required to memorize solubility rules for the upcoming exam. Exam questions will be worded to allow deduction of solubility.

Numerical Examples of Solubility (in grams per 100 grams of water)

  • Sodium Chloride (NaCl): 35.9 g

  • Sodium Nitrate (NaNO_3): 91 g

  • Calcium Acetate (Ca(CH3COO)2): Soluble (Ca^{2+} paired with two acetate ions, CH_3COO^-, a good refresher on Chapter 2 material).

  • Magnesium Sulfate (MgSO_4): Really soluble.

  • Calcium Sulfate (CaSO4): Really insoluble (0.2 g) – remarkably different from MgSO4 despite calcium being directly below magnesium on the periodic table.

  • Lead Chloride (PbCl_2): 4.5 g (on the edge of what is considered soluble)

General Solubility Rules (for familiarity, not memorization)

  • Any ionic compound containing the nitrate ion (NO_3^-$ will be soluble.

  • All salts containing chloride (Cl^-, bromide (Br^-), or iodide (I^-) ions are soluble, except if their partner is Ag^+ (silver), Pb^{2+} (lead), or Hg_2^{2+} (mercury(I)).

  • Nearly all salts of the acetate ion (CH_3COO^-$ are soluble.

  • Most salts of the hydroxide ion (OH^-$ are insoluble (e.g., Mg(OH)2), except for compounds containing Group 1A metal ions or ammonium (NH4^+) ions (e.g., NaOH).

    • Note: Hydroxides with metal cations are generally called bases, but can referred to as salts for their ionization properties.

  • All ionic compounds containing nitrate (NO3^-$), Group 1A metal ions (e.g., Na^+), or ammonium (NH4^+) ions will be soluble.

Precipitation Reactions

  • Definition: A precipitation reaction forms an insoluble salt (a precipitate) by mixing aqueous solutions of two different soluble salts.

  • Example: Mixing silver nitrate (AgNO_3) and sodium chloride (NaCl) solutions.

    • AgNO3 in water exists as separated Ag^+ (aq) and NO3^- (aq) ions.

    • NaCl in water exists as separated Na^+ (aq) and Cl^- (aq) ions.

    • Silver chloride (AgCl) is insoluble in water.

    • When the two solutions are mixed, Ag^+ ions react with Cl^- ions to form solid AgCl precipitate. The Na^+ and NO_3^- ions remain dissolved.

Types of Chemical Equations for Precipitation Reactions

  1. Molecular Equation: Shows the chemical formulas of reactants and products without necessarily reflecting their actual dissolved forms.

    • For soluble ionic compounds in water, "(aq)" is used, implying separation into ions.

    • Example: AgNO3 (aq) + NaCl (aq) \rightarrow AgCl (s) + NaNO3 (aq)

  2. Complete Ionic Equation: Represents all strong electrolytes as dissociated ions in solution.

    • Soluble salts, strong acids, and strong bases are written as separate ions with their charges and "(aq)" designation.

    • Insoluble precipitates are written as their molecular formula with "(s)" designation.

    • Example: Ag^+ (aq) + NO3^- (aq) + Na^+ (aq) + Cl^- (aq) \rightarrow AgCl (s) + Na^+ (aq) + NO3^- (aq)

  3. Net Ionic Equation: Focuses only on the species that undergo a chemical change (i.e., those that form the precipitate).

    • Spectator ions (ions that appear in identical forms on both sides of the complete ionic equation) are omitted.

    • Example: Ag^+ (aq) + Cl^- (aq) \rightarrow AgCl (s)

Spectator Ions

  • Ions in solution not involved in the chemical reaction.

  • They appear in identical forms on both sides of the complete ionic equation.

  • In the AgNO3/NaCl example, Na^+ (aq) and NO3^- (aq) are spectator ions.

Example: Forming Lead Iodide (PbI_2) Precipitate

  • Reactants: Lead nitrate (Pb(NO3)2) (soluble) and potassium iodide (KI) (soluble).

  • Product: PbI_2 (solid, yellow precipitate).

  • Molecular Equation: 2KI (aq) + Pb(NO3)2 (aq) \rightarrow PbI2 (s) + 2KNO3 (aq)

    • Note coefficients for balancing and correct physical states.

  • Complete Ionic Equation: 2K^+ (aq) + 2I^- (aq) + Pb^{2+} (aq) + 2NO3^- (aq) \rightarrow PbI2 (s) + 2K^+ (aq) + 2NO_3^- (aq)

  • Net Ionic Equation: Pb^{2+} (aq) + 2I^- (aq) \rightarrow PbI_2 (s)

  • Spectator Ions: K^+ (aq) and NO_3^- (aq).

Neutralization Reactions (Acid-Base Reactions)

  • Definition: Reactions where acids and bases react with each other.

  • Characteristics:

    • May form a non-electrolyte (like water) or a salt (often a weak electrolyte).

    • Reactants can be soluble or insoluble.

    • Products can sometimes be a gas.

    • Can always be represented by molecular, complete ionic, and net ionic equations.

Case 1: Soluble Substances Producing Water

  • Example: Hydrochloric acid (HCl) and sodium hydroxide (NaOH).

  • $^*$Remember: Strong acids (HCl) and strong bases (NaOH) fully dissociate in water.

  • Molecular Equation: HCl (aq) + NaOH (aq) \rightarrow H_2O (l) + NaCl (aq)

  • Complete Ionic Equation: H^+ (aq) + Cl^- (aq) + Na^+ (aq) + OH^- (aq) \rightarrow H_2O (l) + Na^+ (aq) + Cl^- (aq)

    • HCl exists as H^+ and Cl^- ions.

    • NaOH exists as Na^+ and OH^- ions.

    • H_2O is written as a liquid ("(l)") because it is the solvent; "(aq)" is also acceptable but "(l)" is more precise for the solvent itself.

    • NaCl is written as separated Na^+ and Cl^- ions as it's a soluble salt.

  • Net Ionic Equation: H^+ (aq) + OH^- (aq) \rightarrow H_2O (l)

  • Spectator Ions: Na^+ (aq) and Cl^- (aq).

Case 2: Insoluble Reactants

  • Example: Magnesium hydroxide (Mg(OH)_2) reacting with hydrochloric acid (HCl).

  • $^*$Remember: Mg(OH)_2 is largely insoluble (precipitate) in water, while HCl is a strong acid (soluble).

  • Molecular Equation: Mg(OH)2 (s) + 2HCl (aq) \rightarrow 2H2O (l) + MgCl_2 (aq)

    • Note the coefficient of 2 for HCl to balance the two OH^- ions.

  • Complete Ionic Equation: Mg(OH)2 (s) + 2H^+ (aq) + 2Cl^- (aq) \rightarrow 2H2O (l) + Mg^{2+} (aq) + 2Cl^- (aq)

  • Net Ionic Equation: Mg(OH)2 (s) + 2H^+ (aq) \rightarrow 2H2O (l) + Mg^{2+} (aq)

    • Mg^{2+} is not a spectator ion because it was part of an insoluble solid on the reactant side and appears as a dissolved ion on the product side (its form changed).

  • Spectator Ions: Cl^- (aq).

Case 3: Weak Acids

  • Definition: Weak acids (e.g., HF, acetic acid CH_3COOH) do not fully dissociate in water; most of the acid remains in its molecular form.

  • Writing Ionic Equations for Weak Acids: Represent the acid in its molecular form, not as separated ions.

  • **Example 1: Hydrofluoric Acid (HF) and Sodium Hydroxide (NaOH)

    • HF is a weak acid; NaOH is a strong base.

    • Molecular Equation: HF (aq) + NaOH (aq) \rightarrow H_2O (l) + NaF (aq)

    • Complete Ionic Equation: HF (aq) + Na^+ (aq) + OH^- (aq) \rightarrow H_2O (l) + Na^+ (aq) + F^- (aq)

    • Net Ionic Equation: HF (aq) + OH^- (aq) \rightarrow H_2O (l) + F^- (aq)

  • **Example 2: Acetic Acid (CH_3COOH) and a Strong Base

    • Acetic acid is a weak acid.

    • The only ionizable hydrogen is the one at the end (CH_3COOH$).

    • Net Ionic Equation: CH3COOH (aq) + OH^- (aq) \rightarrow CH3COO^- (aq) + H_2O (l)

Case 4: Polyprotic Acids

  • Definition: Acids that can donate more than one proton (e.g., H2SO4).

  • **Example: Sulfuric Acid (H2SO4) and Potassium Hydroxide (KOH)

    • H2SO4 is a strong polyprotic acid; KOH is a strong base.

    • $^*$Remember: Polyprotic strong acids ionize *completely* for the first proton, but only incompletely for subsequent protons.

    • Molecular Equation: H2SO4 (aq) + 2KOH (aq) \rightarrow 2H2O (l) + K2SO_4 (aq)

      • The two H^+ from H2SO4 require two OH^- from KOH.

      • The spectator ions (K^+ and SO4^{2-}) combine to form soluble K2SO_4$$.