4.C Chemical Reactions and Aqueous Solution

Acids, Bases, and Ionization Review

Focus: exam review of first topics, strong/weak acids and bases, and how to think about electrolytes and solubility.
Key idea: acids are typically written with H+ at the start of their formula; hydrogens not at the start are not ionizable.
Example notes discussed:
- CH₄ contains hydrogens but is not an acid in water because H is not at the start of the formula; those hydrogens are not ionizable in this context.
- HF is a weak acid (H is at the front, not a strong acid).
- HNO₃ (nitric acid) is a strong acid.
- H₂SO₄ (sulfuric acid) is a strong acid.
- CH₄ is not an acid in water, and hydroxide-containing salts (e.g., NaOH) are bases.
Conductivity in solutions:
- Acids in water generally conduct electricity (electrolytes); HF and acetic acid have low conductivity as weak electrolytes.
- NaOH is a strong base and conducts electricity well in solution.
- CH₄ is not water-soluble and would not form an electrolyte solution.
Takeaway: acids, bases, and electrolytes are tied to dissociation in water and the ability to produce ions in solution.

Solubility: soluble vs insoluble salts; the basic predictor

Distinction introduced previously: soluble salts dissolve in water and dissociate; insoluble salts do not dissolve and stay as solids.
Practical simplification for 12 eleven: classify salts as soluble or insoluble; “sparingly soluble” salts are treated as insoluble for practical purposes at high concentrations.
Core goal for today: predict whether a salt formed in a reaction will remain dissolved or precipitate as a solid (solubility prediction).
Two main representations of reactions:
- Ionic (precipitation-focused) vs nonionic (molecular) representations.
- Precipitation reactions are a form of double replacement (ionic switching).
Important note: tables of solubility have patterns and exceptions; rule ordering matters (the most general rule overrides others).

Solubility rules (patterns and exceptions)

Left-hand table (mostly soluble):
- Rule 1: Salts of group 1 cations (Li⁺, Na⁺, K⁺, Cs⁺, etc.) are universally soluble with any anion.
- Rule 2: Nitrates (NO₃⁻) and acetates (C₂H₃O₂⁻) are soluble with any cation; ammonium (NH₄⁺) salts are soluble.
Right-hand table (insoluble by default with notable exceptions):
- Rule 3: Most chlorides, bromides, iodides are soluble, with notable insoluble exceptions (e.g., AgCl, Hg₂Cl₂, PbCl₂).
- Rule 4: Most sulfates are soluble, with key insoluble exceptions (e.g., BaSO₄ is insoluble; other specific cation exceptions exist in some tables).
Other classes (general trend):
- Carbonates (CO₃²⁻), phosphates (PO₄³⁻), sulfides (S²⁻), and hydroxides (OH⁻) are generally insoluble, with exceptions for group 1 cations and NH₄⁺ (and, for some hydroxides, some heavier group 2 cations are slightly soluble depending on table).
Practical approach: memorize the core rules or learn to apply the table by identifying the cation/anions on the list and checking for exceptions; rule 1 supersedes others.
An important reminder: sometimes a salt that is typically insoluble (e.g., carbonates) can be soluble if the cation is group 1 or NH₄⁺; otherwise they are insoluble.

Dissociation and the meaning of dissociation in solution

Dissociation concept: soluble ionic solids dissociate into ions in water; intact formula units are not present as discrete molecules in solution.
Example: Ba(OH)₂(s) dissolves in water and dissociates as
ext{Ba(OH)}_{2(s)}
ightarrow ext{Ba}^{2+}(aq) + 2 ext{OH}^{-}(aq)
In a solution with 5 formula units of Ba(OH)₂, the dissolved species would be 5 Ba^{2+} and 10 OH⁻; there are no intact Ba(OH)₂ formula units left in solution.
This dissociation concept is foundational for understanding precipitation and net ionic equations.

The precipitation problem setup: the “dance” metaphor

Scenario: mix two aqueous salt solutions, AB(aq) and CD(aq). In solution, ions are free and can recombine.
Metaphor: ions are dancers; after mixing, A⁺ can pair with D⁻ and C⁺ can pair with B⁻, forming potentially new salts.
Possible outcomes after mixing:
- If one of the new salts (AD or CB) is insoluble, it precipitates (solid at bottom). Example: AD(s).
- If both new salts are soluble, everything remains in solution (no visible reaction).
Additional nuance: sometimes both potential precipitates could be insoluble, but this is rare and often impractical to observe due to competing solubilities and the likelihood of one precipitate dominating.
The idea of spectator ions: some ions do not participate in the actual chemical change (e.g., K⁺ and NO₃⁻ in a BaSO₄-precipitation scenario) and are effectively spectators.
When a precipitate forms, it is the solid product that comes out of solution; the remaining ions stay in solution.

Predicting precipitation: step-by-step method

Start with two soluble salts in aqueous solution: e.g.,
- 1)
  ext{K}2 ext{SO}4 ext{(aq)}
  ext{Ba(NO}3)2 ext{(aq)}
Step 1: Dissociate each salt into its constituent ions (use the charges and common polyatomic ions):
- ext{K}2 ext{SO}4
  ightarrow 2 ext{K}^+ + ext{SO}_4^{2-}
- ext{Ba(NO}3)2
  ightarrow ext{Ba}^{2+} + 2 ext{NO}_3^-
Step 2: Consider possible new salt combinations formed by pairing cations with anions from the other salt (cross-partnering):
- Potential new salts:
- ext{K}^+ + ext{NO}3^- ightarrow ext{KNO}3 (soluble)
- ext{Ba}^{2+} + ext{SO}4^{2-} ightarrow ext{BaSO}4 (insoluble, precipitate)
Step 3: Check solubility of each potential product using the solubility rules:
- Potassium nitrate (KNO₃) is soluble (Rule 1 & Rule 2).
- Barium sulfate (BaSO₄) is insoluble (solubility exception often listed for sulfates).
Step 4: Write the balanced molecular equation including the precipitate:
ext{K}2 ext{SO}4(aq) + ext{Ba(NO}3)2(aq)
ightarrow ext{BaSO}4(s) + 2 ext{KNO}3(aq)
Step 5: Write the ionic forms to identify spectator ions:
- Complete ionic form:
  egin{aligned} ext{2K}^+(aq) + ext{SO}4^{2-}(aq) + ext{Ba}^{2+}(aq) + 2 ext{NO}3^-(aq)
  ightarrow ext{BaSO}4(s) + 2 ext{K}^+(aq) + 2 ext{NO}3^-(aq)
  \ ext{Spectator ions cancel: } ext{K}^+, ext{NO}3^- ext{ ions remain in solution} \ ext{Net ionic equation: } ext{Ba}^{2+}(aq) + ext{SO}4^{2-}(aq)
  ightarrow ext{BaSO}_4(s)

egin{aligned}
ext{Ba}^{2+}(aq) + ext{SO}4^{2-}(aq) & ightarrow ext{BaSO}4(s)
ag{net ionic}

d
d

\end{aligned}

Step 6: Note about spectator ions: ions that do not participate in the formation of the precipitate are spectators. In this example, K⁺ and NO₃⁻ are spectators; the precipitate is BaSO₄.
Step 7: If no precipitate forms, conclude that no reaction occurs under the given conditions.

Worked example: potassium carbonate and iron(III) chloride

System: K₂CO₃(aq) + FeCl₃(aq)
Step 1: Dissociate into ions:
- ext{K}2 ext{CO}3
  ightarrow 2 ext{K}^+ + ext{CO}_3^{2-}
- ext{FeCl}_3
  ightarrow ext{Fe}^{3+} + 3 ext{Cl}^-
Step 2: Cross-partner to form possible salts:
- ext{K}^+ + ext{Cl}^-
  ightarrow ext{KCl} (soluble)
- ext{Fe}^{3+} + ext{CO}3^{2-} ightarrow ext{Fe}2( ext{CO}3)3(s) (insoluble; precipitate)
Step 3: Write the molecular equation (balanced):
ext{K}2 ext{CO}3(aq) + ext{FeCl}3(aq) ightarrow ext{Fe}2( ext{CO}3)3(s) + 6 ext{KCl}(aq)
Step 4: Complete ionic form:
2 ext{K}^+(aq) + ext{CO}3^{2-}(aq) + 2 ext{Fe}^{3+}(aq) + 6 ext{Cl}^-(aq) ightarrow ext{Fe}2( ext{CO}3)3(s) + 6 ext{K}^+(aq) + 6 ext{Cl}^-(aq)
Step 5: Net ionic equation (spectators canceled):
2 ext{Fe}^{3+}(aq) + 3 ext{CO}3^{2-}(aq) ightarrow ext{Fe}2( ext{CO}3)3(s)
Step 6: Name of the precipitate: iron(III) carbonate. Note on naming: keep the oxidation state in parentheses and write carbonate with no space (Fe₂(CO₃)₃). The discussion emphasizes avoiding spaces in chemical names and formulas.
Step 7: Commentary on balancing and charges: ionic salts must form a neutral compound; sometimes unusual multiples of ions are needed to achieve neutrality (e.g., +6 and −6 charges in this example).

Practice problem walkthrough: aluminum sulfate and barium chloride

System: Al₂(SO₄)₃(aq) + BaCl₂(aq)
Step 1: Dissociate into ions:
- ext{Al}2( ext{SO}4)3 ightarrow 2 ext{Al}^{3+} + 3 ext{SO}4^{2-}
- ext{BaCl}_2
  ightarrow ext{Ba}^{2+} + 2 ext{Cl}^-
Step 2: Cross-partner to form possible salts:
- Potential precipitate: BaSO₄(s) (insoluble)
- Other possible salt: AlCl₃(aq) (soluble, spectator-type salt in this context)
Step 3: Write the balanced molecular equation (and the net ionic):
- Molecular equation:
  ext{Al}2( ext{SO}4)3(aq) + 3 ext{BaCl}2(aq)
  ightarrow 3 ext{BaSO}4(s) + 2 ext{AlCl}3(aq)
- Complete ionic form:
  2 ext{Al}^{3+}(aq) + 3 ext{SO}4^{2-}(aq) + 3 ext{Ba}^{2+}(aq) + 6 ext{Cl}^-(aq) ightarrow 3 ext{BaSO}4(s) + 2 ext{Al}^{3+}(aq) + 6 ext{Cl}^-(aq)
- Net ionic equation (spectators canceled):
  3 ext{Ba}^{2+}(aq) + 3 ext{SO}4^{2-}(aq) ightarrow 3 ext{BaSO}4(s)
Step 4: Identify spectator ions: ions that do not participate in precipitate formation.
- In this reaction, Al^{3+} and Cl^- act as spectator ions (they remain in solution as AlCl₃(aq) or equivalent species).
- The solid precipitate is BaSO₄(s).
Step 5: Answer to the clicker-type prompt:
- The number of spectator ions depends on how you count the species, but the essential spectator species are Al^{3+} and Cl^-, which do not appear in the net ionic equation.
- If asked to name the precipitate, it is barium sulfate. If asked for the spectator salt, it is aluminum chloride/aluminum chloride species in solution.

Key concepts, connections, and implications

Precipitation reactions are governed by solubility rules and the tendency to form insoluble salts from mixing ions.
Spectator ions do not affect the net chemical change; they influence charge balance and overall solution composition but not the precipitation outcome.
The net ionic equation focuses on the ions that actually form the solid precipitate; thermodynamic quantities (enthalpy, free energy) are often determined by these active ions.
Real-world relevance:
- Ocean acidification decreases the solubility of calcium carbonate (CaCO₃), which affects shells and coral structures in marine environments.
- Understanding precipitation underpins purification processes, qualitative analysis, and many industrial separations.
Practical lab considerations:
- Predicting whether a reaction will occur helps in planning experiments, avoiding unnecessary steps, and understanding what will stay in solution versus precipitate.

Formulas, terminology, and quick references (LaTeX)

Dissociation of a soluble salt:
ext{Ba(OH)}_2(s)
ightarrow ext{Ba}^{2+}(aq) + 2 ext{OH}^-(aq)
Solubility-driven precipitation example (molecular):
ext{K}2 ext{SO}4(aq) + ext{Ba(NO}3)2(aq)
ightarrow ext{BaSO}4(s) + 2 ext{KNO}3(aq)
Complete ionic form:
2 ext{K}^+(aq) + ext{SO}4^{2-}(aq) + ext{Ba}^{2+}(aq) + 2 ext{NO}3^-(aq)
ightarrow ext{BaSO}4(s) + 2 ext{K}^+(aq) + 2 ext{NO}3^-(aq)
Net ionic equation:
ext{Ba}^{2+}(aq) + ext{SO}4^{2-}(aq) ightarrow ext{BaSO}4(s)
Example name conventions:
- Iron(III) carbonate: ext{Fe}2( ext{CO}3)_3
- Iron(III) carbonate name: iron(III) carbonate (no spaces in the formula; no extraneous spacing in the name)
General precipitation rule mnemonic: "Dance partners switch; precipitate forms if the new pair is insoluble; spectators stay in solution; if neither new pair is insoluble, no reaction occurs."