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 ($MgSO<em>4$), certain bases like sodium hydroxide ($NaOH$), and certain acids like acetic acid ($CH</em>3COOH$).

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 ($Ca(CH<em>3COO)</em>2$): 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 ($CaSO<em>4$): Really insoluble ($0.2$ g) – remarkably different from $MgSO</em>4$ 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^-$) ions</strong> are soluble, <em>except</em> if their partner is$Ag^+$(silver),$Pb^{2+}$(lead), or$Hg_2^{2+}$(mercury(I)).</p></li><li><p>Nearly all salts of the <strong>acetate ion ($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$).</p><ul><li><p><em>Note</em>: Hydroxides with metal cations are generally called bases, but can referred to as salts for their ionization properties.</p></li></ul></li><li><p>All ionic compounds containing <strong>nitrate ($NO3^-$), Group 1A metal ions (e.g., $Na^+$), or ammonium ($NH</em>4^+$) 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.

  • $AgNO<em>3$ in water exists as separated $Ag^+$ (aq) and $NO</em>3^-$ (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: $AgNO<em>3 (aq) + NaCl (aq) \rightarrow AgCl (s) + NaNO</em>3 (aq)$

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: $Ag^+ (aq) + NO<em>3^- (aq) + Na^+ (aq) + Cl^- (aq) \rightarrow AgCl (s) + Na^+ (aq) + NO</em>3^- (aq)$

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 $AgNO<em>3/NaCl$ example, $Na^+$ (aq) and $NO</em>3^-$ (aq) are spectator ions.

Example: Forming Lead Iodide ($PbI_2$) Precipitate

• Reactants: Lead nitrate ($Pb(NO<em>3)</em>2$) (soluble) and potassium iodide ($KI$) (soluble).

• Product: $PbI_2$ (solid, yellow precipitate).

• Molecular Equation: $2KI (aq) + Pb(NO<em>3)</em>2 (aq) \rightarrow PbI<em>2 (s) + 2KNO</em>3 (aq)$

  • Note coefficients for balancing and correct physical states.

• Complete Ionic Equation: $2K^+ (aq) + 2I^- (aq) + Pb^{2+} (aq) + 2NO<em>3^- (aq) \rightarrow PbI</em>2 (s) + 2K^+ (aq) + 2NO_3^- (aq)$

• 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: $Mg(OH)<em>2 (s) + 2HCl (aq) \rightarrow 2H</em>2O (l) + MgCl_2 (aq)$

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

• Complete Ionic Equation: $Mg(OH)<em>2 (s) + 2H^+ (aq) + 2Cl^- (aq) \rightarrow 2H</em>2O (l) + Mg^{2+} (aq) + 2Cl^- (aq)$

• Net Ionic Equation: $Mg(OH)<em>2 (s) + 2H^+ (aq) \rightarrow 2H</em>2O (l) + Mg^{2+} (aq)$

  • $Mg^{2+}$ 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)$</p></li></ul></li></ul><h5 id="7b44d7ff-f949-48a7-a891-4ea86bc6c02c" data-toc-id="7b44d7ff-f949-48a7-a891-4ea86bc6c02c" collapsed="false" seolevelmigrated="true">Case 4: Polyprotic Acids</h5><ul><li><p><strong>Definition</strong>: Acids that can donate more than one proton (e.g.,$H2SO4$).</p></li><li><p>**Example: Sulfuric Acid ($H2SO4$) and Potassium Hydroxide ($KOH$)</p><ul><li><p>$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)$</p><ul><li><p>The two$H^+$from$H2SO4$require two$OH^-$from$KOH$.</p></li><li><p>The spectator ions ($K^+$and$SO4^{2-}$) combine to form soluble$K2SO_4$$.