Solutions and Electrolytes Quiz

Solutions

• Solutions are homogeneous mixtures of two or more pure substances.
• The solvent is present in the greatest abundance.
• Solutes are the other substances present in the solution.
• An aqueous solution is a solution where water is the solvent.
• Substances dissolve in water through different ways:
  • Ionic compounds dissolve by dissociation, where water surrounds the separate ions.
  • Molecular compounds interact with water but do not dissociate.
  • Some molecular substances react with water when they dissolve.
• The interaction between ions and water molecules involves electrostatic interactions between the ions' charges and the partial charges of the water molecule.
• Chloride ions interact with the partial positive charges of the hydrogen atoms in water, while sodium ions interact with the partial negative charge of the oxygen atom in water.

Electrolytes and Non-Electrolytes

• An electrolyte is a substance that dissociates into ions when dissolved in water.
• A non-electrolyte may dissolve in water but does not dissociate into ions.
• Ionic compounds are strong electrolytes.
• Molecular compounds can be strong electrolytes, weak electrolytes or non-electrolytes.
• Strong acids are strong electrolytes.
• Weak acids and weak bases are weak electrolytes.
• Most molecular compounds are non-electrolytes.

Solubility of Ionic Compounds

• Not all ionic compounds dissolve in water.
• Solubility rules predict whether a combination of ions will dissolve in water.
• General solubility rules:
  • Nitrate, acetate, chloride, bromide, iodide, and sulfate ions are usually soluble with any cation.
  • Chloride, bromide, iodide, and sulfate ions are usually insoluble with silver, mercury, and lead.
  • Sulfate ions with calcium, strontium, and barium are insoluble.
  • Sulfite, carbonate, phosphate, and hydroxide ions are usually insoluble.
  • Exceptions include combinations with ammonium or alkali metals.

Precipitation Reactions

• Mixing two solutions of soluble salts may produce an insoluble salt.
• The insoluble salt precipitates as a solid.
• Precipitation reactions are represented with molecular, complete ionic, and net ionic equations.
• Molecular equations list reactants and products without ionic detail.
• Complete ionic equations show all strong electrolytes dissociated into ions, including spectator ions.
• Net ionic equations show only ions involved in precipitate formation.
• Spectator ions are ions present on both sides of the complete ionic equation, and do not participate in the reaction.

Acids and Bases

• Acids are defined by two main theories:
  • Arrhenius Theory: Acids increase the concentration of protons (H⁺) in water.
  • Brønsted-Lowry Theory: Acids donate protons.
• Bases are defined by two main theories:
  • Arrhenius Theory: Bases increase the concentration of hydroxide ions (OH⁻) in water.
  • Brønsted-Lowry Theory: Bases accept protons.
• Strong acids completely dissociate in water, weak acids only partially dissociate.
• Strong bases completely dissociate in water, weak bases only partially produce hydroxide ions.
• Litmus paper tests for acidity or basicity: Acids turn litmus red, bases turn it blue.

Acid-Base Reactions

• Acid-base reactions involve proton (H⁺) transfer from acid to base.
• Neutralization reaction: Acid reacting with metal hydroxide produces a neutral salt and water.
• Net ionic equation for strong acid/strong base neutralization: H⁺ + OH⁻ → H₂O
• Gas-forming reaction: Carbonates/bicarbonates with acid produce salt, carbon dioxide (CO₂), and water.

Oxidation-Reduction (Redox) Reactions

• Oxidation is electron loss.
• Reduction is electron gain.
• Oxidation numbers track electron transfer.
• Rules for assigning oxidation numbers:
  • Elements in their elemental form have a 0 oxidation number.
  • Monatomic ion's oxidation number is its charge.
  • Nonmetals tend to have negative oxidation numbers (oxygen usually -2, except in peroxides).
  • Hydrogen usually +1 (except when bonded to metals).
  • Fluorine always -1.
  • Sum of oxidation numbers in a neutral compound is zero.
  • Sum of oxidation numbers in a polyatomic ion equals the charge of the ion.
• Displacement reactions are a type of redox reaction, where one ion oxidizes another element.
• Metals above hydrogen in the activity series more readily react with acids, producing hydrogen gas.

Concentration

• Molarity (M) measures moles of solute per liter of solution.
• Molarity = moles of solute / volume of solution (in liters)
• Solutions of known molarity are prepared by dissolving a calculated mass of solute in a volumetric flask to a specific volume.
• Dilution reduces concentration by adding more solvent.
• Dilution equation: M₁V₁ = M₂V₂ (initial concentration x initial volume = final concentration x final volume)
• Stoichiometric calculations can be used with molarity to determine reactants/products.

Titration

• Titration uses a solution of known concentration (standard solution) to find an unknown solution's concentration.
• Equivalence point: Moles of titrant = moles of analyte.
• Unknown concentration is calculated from the known solution's volume and concentration, stoichiometry, and unknown solution volume.