Reactions in Aqueous Solution - Vocabulary Flashcards

Solutions and Solvation

• Solution Definitions:

  • Solution: A homogeneous mixture composed of two or more pure substances.

  • Solvent: The component of a solution present in the greatest abundance.

  • Solute: Any substance dissolved in the solvent; all non-solvent components in the mixture.

  • Aqueous Solution: A solution in which water serves as the solvent.

• Solvation Mechanism:

  • Solvation: The universal process by which solute particles dissolve, defined as the surrounding of solute particles by solvent molecules.

  • Ionic Compound Dissolution: Dissolve via dissociation, where water molecules surround individual, separated cations and anions.

  • Molecular Compound Dissolution: Disperse evenly throughout water while most molecules remain intact as neutral units; a subset of molecular substances chemically react with water to form ions upon dissolving.

Electrolytes and Dissociation Behavior

• Electrolyte vs. Nonelectrolyte:

  • Electrolyte: A substance that dissociates into ions when dissolved in water, producing a solution that conducts electricity (e.g., sodium chloride, $\text{NaCl}(s)$, and acetic acid, $\text{CH}_3\text{COOH}(aq)$).

    • $\text{NaCl}(s)$ dissociates into $\text{Na}^+(aq)$ and $\text{Cl}^-(aq)$.

    • $\text{CH}_3\text{COOH}(aq)$ dissociates into $\text{CH}_3\text{COO}^-(aq)$ and $\text{H}^+(aq)$.

  • Nonelectrolyte: A substance that may dissolve in water but does not dissociate into ions upon doing so. Nonelectrolyte solutions do not conduct electricity (e.g., sugars like glucose and sucrose).

    • Glucose dissolution: $\text{C}_6\text{H}_{12}\text{O}_6(s) \rightarrow \text{C}_6\text{H}_{12}\text{O}_6(aq)$

    • Sucrose dissolution: $\text{C}_{12}\text{H}_{22}\text{O}_{11}(s) \rightarrow \text{C}_{12}\text{H}_{22}\text{O}_{11}(aq)$

• Classification of Electrolytic Strength:

  • Strong Electrolyte: Dissociates completely into ions when dissolved in water. Equations feature a single forward arrow ($\rightarrow$). The solution acts as a strong electrical conductor.

    • Dissociation equation: $\text{HCl}(aq) \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)$

  • Weak Electrolyte: Dissociates only partially into ions when dissolved in water. Dissociation equations reflect chemical equilibrium using a double reversible arrow ($\rightleftharpoons$). The solution acts as a weak electrical conductor.

    • Equilibrium equation: $\text{CH}_3\text{COOH}(aq) \rightleftharpoons \text{H}^+(aq) + \text{CH}_3\text{COO}^-(aq)$

  • Nonelectrolyte: Does not dissociate into ions when dissolved in water and yields a non-conducting solution.

Solubility Guidelines for Ionic Compounds

• Precipitation Fundamentals:

  • Precipitation Reaction: Occurs when two aqueous solutions containing soluble salts are mixed, reacting to form an insoluble ionic solid.

  • Precipitate: The insoluble solid compound that forms and separates from the solution during a precipitation reaction.

• Solubility Guidelines in Water (Table 4.1):

  • Soluble Ionic Compounds:

    • Nitrate ($\text{NO}_3^-$): All nitrate compounds are soluble; no exceptions.

    • Acetate ($\text{CH}_3\text{COO}^-$): All acetate compounds are soluble; no exceptions.

    • Chloride ($\text{Cl}^-$): Soluble, except compounds containing $\text{Ag}^+$, $\text{Hg}_2^{2+}$, and $\text{Pb}^{2+}$.

    • Bromide ($\text{Br}^-$): Soluble, except compounds containing $\text{Ag}^+$, $\text{Hg}_2^{2+}$, and $\text{Pb}^{2+}$.

    • Iodide ($\text{I}^-$): Soluble, except compounds containing $\text{Ag}^+$, $\text{Hg}_2^{2+}$, and $\text{Pb}^{2+}$.

    • Sulfate ($\text{SO}_4^{2-}$): Soluble, except compounds containing $\text{Sr}^{2+}$, $\text{Ba}^{2+}$, $\text{Hg}_2^{2+}$, and $\text{Pb}^{2+}$.

  • Insoluble Ionic Compounds:

    • Sulfide ($\text{S}^{2-}$): Insoluble, except compounds containing $\text{NH}_4^+$, alkali metal cations, $\text{Ca}^{2+}$, $\text{Sr}^{2+}$, and $\text{Ba}^{2+}$.

    • Carbonate ($\text{CO}_3^{2-}$): Insoluble, except compounds containing $\text{NH}_4^+$ and alkali metal cations.

    • Phosphate ($\text{PO}_4^{3-}$): Insoluble, except compounds containing $\text{NH}_4^+$ and alkali metal cations.

    • Hydroxide ($\text{OH}^-$): Insoluble, except compounds containing $\text{NH}_4^+$, alkali metal cations, $\text{Ca}^{2+}$, $\text{Sr}^{2+}$, and $\text{Ba}^{2+}$.

Metathesis Reactions and Writing Ionic Equations

• Metathesis (Exchange) Reaction Principles:

  • Etymology: Derived from Greek meaning "to transpose".

  • General reaction scheme: $AX + BY \rightarrow AY + BX$

  • Mechanism: The anions ($X^-$ and $Y^-$) in the reactant compounds exchange or transpose between cations ($A^+$ and $B^+$).

  • Example reaction: $\text{Mg(NO}_3)_2(aq) + 2\text{NaOH}(aq) \rightarrow \text{Mg(OH)}_2(s) + 2\text{NaNO}_3(aq)$, where $X = \text{NO}_3^-$ and $Y = \text{OH}^-$.

• Predicting Precipitate Formation:

  1. Note all individual ions present in the reactant solutions.

  2. Write all possible cation-anion combinations.

  3. Evaluate each combination against the solubility rules to determine if any product is insoluble.

  4. Example analysis for mixing $\text{Mg(NO}_3)_2(aq)$ and $\text{NaOH}(aq)$:

     • Reactant Cations: $\text{Mg}^{2+}$, $\text{Na}^+$

     • Reactant Anions: $\text{NO}_3^-$, $\text{OH}^-$

     • Product Combinations: $\text{NaNO}_3$ (soluble) and $\text{Mg(OH)}_2$ (insoluble solid precipitate).

• Steps to Complete and Balance Metathesis Equations:

  1. Identify constituent reactant ions from chemical formulas.

  2. Write product chemical formulas by pairing the cation of one reactant with the anion of the other, balancing ionic charges for proper subscripts.

  3. Consult solubility guidelines to determine if any product is insoluble.

  4. Balance the overall chemical equation.

• Procedure for Writing Net Ionic Equations:

  1. Molecular Equation: Write the balanced metathesis equation without depicting ionic dissociation, specifying state symbols ($(aq)$, $(s)$, $(l)$, $(g)$).

     • Example: $\text{Pb(NO}_3)_2(aq) + 2\text{KCl}(aq) \rightarrow \text{PbCl}_2(s) + 2\text{KNO}_3(aq)$

  2. Complete Ionic Equation: Rewrite the equation showing all soluble strong electrolytes (strong acids, strong bases, soluble ionic salts) dissociated into their individual ions. Only strong electrolytes dissolved in aqueous solution are split into ions.

     • Example: $\text{Pb}^{2+}(aq) + 2\text{NO}_3^-(aq) + 2\text{K}^+(aq) + 2\text{I}^-(aq) \rightarrow \text{PbI}_2(s) + 2\text{K}^+(aq) + 2\text{NO}_3^-(aq)$

  3. Net Ionic Equation: Identify and eliminate spectator ions—ions present on both sides of the reaction that remain unchanged during the reaction.

     • Spectator ions in example: $\text{K}^+$ and $\text{NO}_3^-$

     • Net ionic equation: $\text{Pb}^{2+}(aq) + 2\text{I}^-(aq) \rightarrow \text{PbI}_2(s)$

Acids, Bases, and Neutralization Reactions

• Definitions of Acids and Bases:

  • Acid: A substance that ionizes in aqueous solution to yield hydrogen ions ($\text{H}^+$); defined as a proton donor.

  • Base: A substance that accepts or reacts with $\text{H}^+$ ions and increases hydroxide ion ($\text{OH}^-$) concentration when dissolved in water. A compound does not require an $\text{OH}^-$ group in its formula to function as a base (e.g., ammonia, $\text{NH}_3$).

• Acid and Base Dissociation Strengths:

  • Strong Acid: Completely ionizes in aqueous solution.

  • Weak Acid: Only partially ionizes in aqueous solution.

  • Strong Base: Dissociates completely into metal cations and hydroxide anions in water.

  • Weak Base: Only partially reacts with water to generate hydroxide anions.

• Common Strong Acids and Strong Bases (Table 4.2):

  • Strong Acids:

    • Hydrochloric acid, $\text{HCl}$

    • Hydrobromic acid, $\text{HBr}$

    • Hydroiodic acid, $\text{HI}$

    • Chloric acid, $\text{HClO}_3$

    • Perchloric acid, $\text{HClO}_4$

    • Nitric acid, $\text{HNO}_3$

    • Sulfuric acid (first proton dissociation only), $\text{H}_2\text{SO}_4$

  • Strong Bases:

    • Group 1A metal hydroxides: $\text{LiOH}$, $\text{NaOH}$, $\text{KOH}$, $\text{RbOH}$, $\text{CsOH}$

    • Heavy Group 2A metal hydroxides: $\text{Ca(OH)}_2$, $\text{Sr(OH)}_2$, $\text{Ba(OH)}_2$

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• Systematic Classification of Electrolytic Behavior (Table 4.3):

  1. Determine if the solute is ionic or molecular.

  2. If Ionic: All soluble ionic compounds are Strong Electrolytes. None are weak electrolytes or nonelectrolytes.

  3. If Molecular:

     • Acid test: Starts with H or ends in $\text{COOH}$. If listed as a strong acid, it is a Strong Electrolyte. If not on the strong acid list, it is a weak acid and thus a Weak Electrolyte.

     • Base test: Strong bases are Strong Electrolytes. Ammonia ($\text{NH}_3$) is a weak base and thus a Weak Electrolyte.

     • All other molecular compounds are Nonelectrolytes.

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• Neutralization Reactions:

  • Definition: A reaction taking place between an acid and a base.

  • Reaction Products: When the reacting base is a metal hydroxide, the products are water and a salt (an ionic compound).

  • Neutralization Equations (e.g., $\text{HCl}$ + $\text{NaOH}$):

    • Molecular Equation: $\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{H}_2\text{O}(l) + \text{NaCl}(aq)$

    • Complete Ionic Equation: $\text{H}^+(aq) + \text{Cl}^-(aq) + \text{Na}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l) + \text{Na}^+(aq) + \text{Cl}^-(aq)$

    • Net Ionic Equation: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$

    • Special Rule: If a weak electrolyte participates in a neutralization reaction, it is written in its complete molecular form rather than separated into ions.

• Neutralization Reactions with Gas Formation:

  • Carbonates and Bicarbonates: React with acids to form a salt, carbon dioxide gas ($\text{CO}_2$), and liquid water.

    • $\text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{CO}_2(g) + \text{H}_2\text{O}(l)$

    • $\text{NaHCO}_3(aq) + \text{HBr}(aq) \rightarrow \text{NaBr}(aq) + \text{CO}_2(g) + \text{H}_2\text{O}(l)$

  • Sulfides: React with acids to form a salt and hydrogen sulfide gas ($\text{H}_2\text{S}$), which is poisonous and has a distinct rotten-egg odor.

    • $\text{Na}_2\text{S}(aq) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{Na}_2\text{SO}_4(aq) + \text{H}_2\text{S}(g)$

• Commercial Antacids and Active Ingredients (Table 4.4):

  • Alka-Seltzer®: $\text{NaHCO}_3$

  • Amphojel®: $\text{Al(OH)}_3$

  • Di-Gel®: $\text{Mg(OH)}_2$ and $\text{CaCO}_3$

  • Milk of Magnesia: $\text{Mg(OH)}_2$

  • Maalox®: $\text{Mg(OH)}_2$ and $\text{Al(OH)}_3$

  • Mylanta®: $\text{Mg(OH)}_2$ and $\text{Al(OH)}_3$

  • Rolaids®: $\text{Mg(OH)}_2$ and $\text{CaCO}_3$

  • Tums®: $\text{CaCO}_3$

Oxidation-Reduction Reactions

• Fundamental Redox Concepts:

  • Oxidation: Loss of electrons.

  • Reduction: Gain of electrons.

  • Oxidation and reduction must always occur simultaneously; oxidation cannot happen without reduction.

• Oxidation Numbers:

  • An assigned value for each atom in a neutral compound or charged entity used for electron accounting ("bookkeeping").

  • Oxidation numbers do not represent actual ionic charges on atoms.

• Rules for Assigning Oxidation Numbers:

  1. Elemental State: Atoms in elemental form have an oxidation number of $0$ (e.g., $\text{Na}(s)$, $\text{Fe}(s)$, $\text{H}_2(g)$, $\text{O}_2(g)$, $\text{N}_2(g)$, $\text{Cl}_2(g)$).

  2. Monatomic Ions: Oxidation number equals the ionic charge.

     • Charge $+1$: H, Li, Na, K

     • Charge $+2$: Be, Mg, Ca

     • Charge $+3$: B, Al, Ga

     • Charge $+4$ (or $-2$): C, Si, Ge

     • Charge $-1$: F, Cl, Br, I

     • Charge $-2$: O

     • Charge $-3$: N

  3. Nonmetals: Generally have negative oxidation numbers, though positive values occur:

     • Oxygen: Typically $-2$, except in the peroxide ion ($\text{O}_2^{2-}$) where it is $-1$.

     • Hydrogen: $+1$ when bonded to nonmetals; $-1$ when bonded to metals.

     • Fluorine: Always $-1$.

     • Other Halogens: Typically $-1$, but positive ($+1$, $+3$, $+5$, $+7$) when combined with oxygen in oxyanions.

  4. Sum of Oxidation Numbers:

     • The sum of oxidation numbers in a neutral compound is $0$ (e.g., $\text{H}_2\text{O}$, $\text{AlCl}_3$, $\text{Mg}_3\text{N}_2$).

     • The sum of oxidation numbers in a polyatomic ion equals the overall charge on the ion (e.g., $\text{Cr}_2\text{O}_7^{2-}$, $[\text{Al(OH)}_4]^-$, $\text{S}_2\text{O}_3^{2-}$, $\text{SO}_4^{2-}$).

     • Every atom must be counted according to its chemical subscript.

• Displacement Reactions:

  • Reactions where an ion in solution oxidizes an elemental metal.

  • Example: $\text{H}^+(aq)$ oxidizes $\text{Mg}(s)$, displacing $\text{H}^+$ ions in solution to yield $\text{H}_2(g)$.

• Activity Series of Metals (Table 4.5):

  • Predicts reaction feasibility based on metal reactivity.

  • Metals positioned above hydrogen in the activity series react with acids to produce hydrogen gas ($\text{H}_2$). Metals below hydrogen will not react with acids.

  • A more reactive metal (higher on the activity series) is oxidized to a cation, displacing the lower metal ion which is reduced to its elemental solid state (e.g., copper ($\text{Cu}$) is above silver ($\text{Ag}$); $\text{Cu}(s)$ oxidizes to $\text{Cu}^{2+}(aq)$ while $\text{Ag}^+(aq)$ reduces to $\text{Ag}(s)$).

Solution Molarity and Dilution

• Concentration Concept:

  • Concentration describes the quantity of solute dissolved in a given amount of solution.

• Molarity Formula:

  • Molarity ($M$) is defined as the number of moles of solute per liter of solution:         $\text{Molarity } (M) = \frac{\text{moles of solute}}{\text{volume of solution in liters}}$

  • Serves as a conversion factor between moles of solute and volume of solution in liters.

• Procedure for Preparing Solutions:

  1. Weigh out the calculated mass (and corresponding moles) of solute.

  2. Place solute into a volumetric flask.

  3. Add solvent until the solution level reaches the volumetric mark on the flask neck.

• Dilution Principles and Calculations:

  • Dilution is performed by adding only solvent to a solution.

  • Dilution decreases concentration while leaving the total number of moles of solute unchanged.

  • Dilution Equation:         $M_c \times V_c = M_d \times V_d$         where $M_c$ and $M_d$ are the molarities of concentrated and dilute solutions, and $V_c$ and $V_d$ are their respective volumes.

  • Relationship between concentration, volume, and moles:         $\text{Moles of solute} = \text{Molarity } (M) \times \text{Volume in liters } (V)$         $\frac{\text{mol}}{\text{L}} \times \text{L} = \text{mol}$

Solution Stoichiometry and Titrations

• Solution Stoichiometry:

  • Applies balanced chemical equations alongside solution molarity and volume values to calculate quantitative reactant and product relationships.

• Titration Technique:

  • An analytical technique used to determine the unknown concentration of a solute in a solution.

  • Standard Solution: A reagent solution of precisely known concentration used to react with the solution of unknown concentration.

  • Equivalence Point: The point in a titration where stoichiometric equivalents of reactants have completely reacted.

  • End Point: The experimentally observed point signaling reaction completion, typically made visible by a color change of an indicator.