Solutions and Electrolytes

Solutions and Molecular Compounds

• Molecular compounds typically remain intact when dissolved.
• Example: Oxygen molecules dissolve in water, allowing fish to breathe.
• CO(2) reacts with water, forming carbonic acid (HCO(3^-\
  ) and bicarbonate (HCO(_3^-")).

Carbonation Example

• CO(_2) increases solubility under pressure, such as in carbonated drinks.
• Reaction:
• CO(2) + H(2)O ⇌ H(2)CO(3) (carbonic acid)
• Carbonic acid dissociates into bicarbonate and hydrogen ions.

Mentos and Diet Coke

• Mentos acts as a catalyst for the rapid release of CO(_2) in Diet Coke.
• The exact mechanism of this reaction is not fully understood but leads to explosive eruptions.

Molarity Introduction

• Molarity is essential for calculations in solution chemistry.
• The chapter will cover qualitative aspects of solutions and chemical reactions, emphasizing deeper understanding rather than just numerical values.

Chemical Bonding and Compounds

• Types of compounds: covalent and ionic.

• Ionic compounds involve electron transfer (e.g., NaCl).

• Ionic compounds consist of cations (e.g., Na(^+)) and anions (e.g., Cl(^-)).

• Covalent compounds involve shared electrons (e.g., Cl(_2)).

• Show a spectrum of polarity – between ionic and covalent.

Polarity

• Defined as the distribution of positive and negative charges within a molecule.
• Water (H(_2)O) is highly polar and exhibits unique properties.
• Comparison: Alkanes (nonpolar) vs. Alcohols (polar).
• Example: Ethanol is polar and soluble in water.

Like Dissolves Like

• Polar solvents like water dissolve polar compounds.
• Example: Sodium chloride dissolves in water due to ion-dipole interactions.
• Opposite: Oil separates from water due to nonpolar nature.

Electrolytes

• Strong Electrolyte: Fully dissociates in solution (e.g., NaCl).
• Weak Electrolyte: Partial dissociation (e.g., acetic acid).
• Non-electrolyte: Does not dissociate (e.g., sugar).
• Conductivity in solutions depends on the presence of charged ions.
• Pure water conducts poorly unless electrolytes are present.

Precipitation Reactions

• Occur when two aqueous solutions yield an insoluble solid (precipitate).
• Example: Mixing silver nitrate and sodium chloride forms silver chloride solid.
• Double displacement reactions lead to precipitate formation:
• Example Reaction:
  • AgNO(3)(aq) + NaCl(aq) → AgCl(s) + NaNO(3)(aq)

Ionic Equations

• Complete Ionic Equation: Break down all strong electrolytes into their ions.
• Net Ionic Equation: Focus on ions that participate in forming the precipitate.
• Example: Ag(^+)(aq) + Cl(^-)(aq) → AgCl(s) (precipitate)

Solubility Rules

• General rules guide predictions of solubility for ionic compounds.
• Example Solubility:
  • Most nitrates are soluble.
  • Chlorides are generally soluble except for Ag(^+), Pb(^{2+}), and Hg(^{2+}) compounds.

Applying Solubility Rules

• Use solubility tables to predict outcomes of mixing solutions.
• Example: Determining precipitate formation with NaI and Pb(NO₃)₂ involves checking solubilities of possible products (e.g., PbI₂) and spectator ions (unchanged ions).