Chemical Reactions and Aqueous Solutions

Chemical Equations

• Information in a Chemical Equation

  • Chemical equations serve as symbolic representations of chemical reactions where reactants rearrange their bonding to form products.
  • Identity of substances is provided by chemical formulas.
  • Proportions of the reaction are provided by coefficients.
  • Example Reaction: Two hydrogen molecules react with one oxygen molecule to form two water molecules.
    • $2 H_2 + O_2 \rightarrow 2 H_2O$
    • Reactants: $2 H_2$ and $O_2$
    • Product: $2 H_2O$

• Physical States of Reactants and Products

  • Information in parentheses following a formula identifies the physical state:
    • Solid: $(s)$
    • Liquid: $(l)$
    • Gas: $(g)$
    • Aqueous solution (dissolved in water): $(aq)$
  • Example: $MgI_2(aq) + 2 AgNO_3(aq) \rightarrow 2 AgI(s) + Mg(NO_3)_2(aq)$
    • Interpretation: $1\,mol$ of aqueous magnesium iodide reacts with $2\,mol$ of aqueous silver nitrate to produce $2\,mol$ of solid silver iodide and $1\,mol$ of aqueous magnesium nitrate.

• Reaction Conditions

  • Conditions required for a reaction (like specific temperatures or catalysts) are written above or below the reaction arrow.
  • Example: The decomposition of magnesium carbonate.
    • $MgCO_3(s) \xrightarrow{\text{heat}} MgO(s) + CO_2(g)$
    • Heat is mandatory; without it, the reaction will not occur.

• Balancing Chemical Equations

  • Law of Conservation of Mass: All atoms present at the start of a reaction must be present at the end. The number of atoms for each element must be equal on both the reactant (left) and product (right) sides.
  • Correct Procedure: Only change the coefficients (the numbers in front of formulas). Never change the subscripts within a formula, as changing biological subscripts change the identity of the chemical compound.
  • Balancing Strategy:
    1. If polyatomic ions (e.g., $SO_4^{2-}$) appear on both sides of the equation, balance them as a single unit rather than as individual elements.
    2. Balance elements that appear in only one reactant and one product first.
    3. Balance elements that appear in multiple reactants or products last.
    4. Verify the count of each atom and ensure coefficients are in the smallest whole-number ratio.
    5. If a fractional coefficient is used, multiply the entire equation by the denominator to achieve whole numbers.

• Example 4.1: Barium Hydroxide and Hydrobromic Acid

  • Unbalanced: $Ba(OH)_2 + HBr \rightarrow BaBr_2 + H_2O$
  • Balanced: $Ba(OH)_2 + 2 HBr \rightarrow BaBr_2 + 2 H_2O$

• Example 4.3: Potassium Iodate and Potassium Iodide

  • Reaction: $KIO_3 + KI + HCl \rightarrow I_2 + H_2O + KCl$
  • Balanced: $KIO_3 + 5 KI + 6 HCl \rightarrow 3 I_2 + 3 H_2O + 6 KCl$

• Example 4.4: Sodium Hydroxide and Phosphoric Acid

  • Unbalanced: $NaOH + H_3PO_4 \rightarrow Na_2HPO_4 + H_2O$
  • Balanced: $2 NaOH + H_3PO_4 \rightarrow Na_2HPO_4 + 2 H_2O$

Types of Chemical Reactions

• Basic Reaction Patterns

  • Synthesis (Combination): Simple reactants combine to form one complex product. $A + B \rightarrow AB$
    • Example: Rusting of iron metal ($4 Fe(s) + 3 O_2(g) \rightarrow 2 Fe_2O_3(s)$).
  • Decomposition: A single reactant breaks down into less complex products. $AB \rightarrow A + B$
    • Example: Electrolysis of water ($2 H_2O(l) \xrightarrow{\text{electricity}} 2 H_2(g) + O_2(g)$).
    • Example: Decomposition of potassium chlorate ($2 KClO_3(s) \rightarrow 2 KCl(s) + 3 O_2(g)$).
  • Single-Replacement (Displacement): An element reacts with a compound and displaces another element. $A + BC \rightarrow AC + B$
    • Example: Zinc reacting with hydrochloric acid ($Zn(s) + 2 HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$).
  • Double-Replacement (Metathesis): Two ionic compounds exchange ions to form two new compounds. $AB + CD \rightarrow AD + CB$
    • Example: $2 KI(aq) + Pb(NO_3)_2(aq) \rightarrow PbI_2(s) + 2 KNO_3(aq)$.
  • Combustion: Rapid combination of a substance with oxygen, typically producing heat.
    • Hydrocarbon Pattern: $C_xH_y + O_2 \rightarrow CO_2 + H_2O$
    • Example: Propane combustion ($C_3H_8(g) + 5 O_2(g) \rightarrow 3 CO_2(g) + 4 H_2O(l)$).
  • Acid-Base: A specific type of double-replacement where an acid reacting with a base produces a salt and water.
    • General: $\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$
    • Example: $HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$.

• Driving Forces for Reactions

  • Reactions occur spontaneously toward stable, lower-energy products based on enthalpy and entropy.
  • Precipitation: The driving force for double-replacement reactions that form an insoluble solid ionic compound.
  • Neutralization: The driving force for acid-base reactions forming water and salt.
  • Oxidation-Reduction (Redox): The driving force for synthesis, decomposition, and single-replacement reactions involving electron transfer.

Compounds in Aqueous Solution

• Aqueous Behavior

  • Soluble: A compound that readily dissolves in water.
  • Insoluble: A compound that remains a solid in water.
  • Mobility: Dissolved compounds are more mobile than solids and more likely to collide and react.

• Dissociation and Electrolytes

  • Hydrated Ions: When ionic compounds dissolve, they dissociate into individual ions surrounded by water molecules. $NaCl(s) \xrightarrow{H_2O} Na^+(aq) + Cl^-(aq)$.
  • Electrolytes: Substances that conduct electricity when dissolved in water due to mobile ions.
  • Strong Electrolytes: Dissociate $100\%$ in water (e.g., ionic salts like $NaCl$, strong acids, and strong bases).
  • Weak Electrolytes: Ionize only partially; solutions consist mostly of molecules with few ions (e.g., weak acids and weak bases like $NH_3$).
  • Nonelectrolytes: Dissolve as whole molecules and do not conduct electricity (e.g., sugars like $C_6H_{12}O_6$).

• Acids and Bases in Water

  • Strong Acids: Ionize $100\%$ in water.
    • $HCl$ (Hydrochloric acid)
    • $HBr$ (Hydrobromic acid)
    • $HI$ (Hydroiodic acid)
    • $HNO_3$ (Nitric acid)
    • $HClO_4$ (Perchloric acid)
    • $HClO_3$ (Chloric acid)
    • $H_2SO_4$ (Sulfuric acid)
  • Weak Acids: Any acid not on the strong acid list (e.g., $HNO_2$, $H_3PO_4$, $HF$, $HC_2H_3O_2$).
  • Strong Bases: Ionic compounds containing hydroxide ($OH^-$) that dissociate $100\%$ (e.g., $NaOH$, $KOH$, $Ba(OH)_2$).
  • Weak Bases: Molecular compounds that produce hydroxide ions only to a small extent (e.g., ammonia, $NH_3$).

Precipitation Reactions

• Predicting Solubility

  • Solubility levels: Soluble (>$1\,g$ solute/$100\,g$ water), Slightly Soluble ($0.1\,g$ to $1\,g$), Insoluble (<$0.1\,g$).
  • Solubility Guidelines:
    1. Group 1 elements ($Li^+$, $Na^+$, $K^+$, $Rb^+$, $Cs^+$, $Fr^+$) and Ammonium ($NH_4^+$) salts are soluble.
    2. Nitrates ($NO_3^-$), Chlorates ($ClO_3^-$), Perchlorates ($ClO_4^-$), and Acetates ($C_2H_3O_2^-$) are soluble.
    3. Chlorides ($Cl^-$), Bromides ($Br^-$), and Iodides ($I^-$) are soluble, except for those joined with $Ag^+$, $Pb^{2+}$, or $Hg_2^{2+}$.
    4. Carbonates ($CO_3^{2-}$), Sulfites ($SO_3^{2-}$), Phosphates ($PO_4^{3-}$), and Chromates ($CrO_4^{2-}$) are insoluble, except with guideline 1 cations.
    5. Hydroxides ($OH^-$) and Sulfides ($S^{2-}$) are insoluble, except with guideline 1 cations and $Ba^{2+}$.
    6. Silver ($Ag^+$), mercury ($Hg_2^{2+}$), and lead ($Pb^{2+}$) salts are generally insoluble (unless paired with guideline 2 anions).
    7. Sulfates ($SO_4^{2-}$) are soluble, except with $Ca^{2+}$, $Sr^{2+}$, $Ba^{2+}$, $Ag^+$, $Pb^{2+}$, or $Hg_2^{2+}$.

• Types of Equations for Aqueous Reactions

  • Total (Overall) Equation: Shows complete neutral formulas. $AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$.
  • Ionic Equation: Shows all strong electrolytes as dissociated ions. $Ag^+(aq) + NO_3^-(aq) + Na^+(aq) + Cl^-(aq) \rightarrow AgCl(s) + Na^+(aq) + NO_3^-(aq)$.
  • Net Ionic Equation: Excludes spectator ions (ions that appear unchanged on both sides) to show only reacting species. $Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)$.

Acid-Base Reactions

• Product Prediction

  • Hydrogen from the acid ($H^+$) combines with hydroxide from the base ($OH^-$) to form water ($H_2O(l)$).
  • The remaining cation (from base) and anion (from acid) form a salt.
  • Example: $2 HClO_4(aq) + Ba(OH)_2(aq) \rightarrow Ba(ClO_4)_2(aq) + 2 H_2O(l)$.

• Net Ionic Equations for Acid-Base

  • Strong Acid + Strong Base: The net ionic equation is always $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$.
  • Weak Acid + Strong Base: The weak acid is written as a whole molecule because it does not ionize significantly.
    • Example (Weak acid $HF$): $HF(aq) + OH^-(aq) \rightarrow F^-(aq) + H_2O(l)$.
    • Example (Weak acid $HC_2H_3O_2$): $HC_2H_3O_2(aq) + OH^-(aq) \rightarrow C_2H_3O_2^-(aq) + H_2O(l)$.

Oxidation States and Redox Reactions

• Rules for Assigning Oxidation States (Oxidation Numbers)

  1. A neutral element not in a compound is $0$ (e.g., $Mg$, $O_2$).
  2. Monoatomic ions equal their charge (e.g., $P^{3-}$ is $-3$; $Fe^{3+}$ in $FeCl_3$ is $+3$).
  3. Sum of oxidation states for a neutral compound is $0$; for a polyatomic ion, it equals the ion charge.
  4. Oxygen is usually $-2$ in compounds.
  5. Hydrogen is usually $+1$ in compounds.
  • Example Calculation: Find Phosphorus in $K_3PO_4$.
    • $3(+1) + x + 4(-2) = 0$
    • $3 + x - 8 = 0 \rightarrow x = +5$.

• Redox Terminology

  • Oxidation: An increase in oxidation state; loss of electrons.
  • Reduction: A decrease (reduction) in oxidation state; gain of electrons.
  • Oxidizing Agent: The reactant that accepts electrons (it gets reduced).
  • Reducing Agent: The reactant that provides electrons (it gets oxidized).
  • Example: Synthesis of Aluminum Chloride ($2 Al + 3 Cl_2 \rightarrow 2 AlCl_3$
    • $Al$ goes from $0$ to $+3$ (Oxidized, Reducing Agent).
    • $Cl$ goes from $0$ to $-1$ (Reduced, Oxidizing Agent).

Predicting the Products of Redox Reactions

• Synthesis and Decomposition

  • Often involve redox. Products can be predicted using characteristic monatomic charges for Groups 1, 2, 16, 17, and Aluminum.
  • Examples:
    • $2 Rb(s) + F_2(g) \rightarrow 2 RbF(s)$
    • $3 Ca(s) + N_2(g) \rightarrow Ca_3N_2(s)$

• Single-Replacement and the Activity Series

  • Active metals (strongest reducing agents) give up electrons readily to ions of less active metals.
  • The Activity Series (Most active to least active):
    1. Groups 1-2 metals ($Li$, $K$, $Ca$, $Na$, etc.)
    2. $Al$
    3. $Mn$
    4. $Zn$
    5. $Cr$
    6. $Fe$
    7. $Ni$
    8. $Sn$
    9. $Pb$
    10. $H$ (Standard for comparison)
    11. $Cu$
    12. $Ag$
    13. $Au$
  • Predicting Reaction:
    • A reaction occurs only if the neutral metal is higher on the activity series than the ion it is replacing.
    • $Li(s) + Cu^{2+}(aq) \rightarrow 2 Li^+(aq) + Cu(s)$ (Occurs because $Li$ is above $Cu$).
    • $Cu(s) + Li^+(aq) \rightarrow \text{No Reaction}$ (Because $Cu$ is below $Li$).

• Reactions of Metals with Acids

  • Neutral metals located above hydrogen (H) in the activity series react with aqueous acids to form $H_2(g)$.
  • Metals below H (like $Cu$, $Ag$, $Au$) do not react with acids.
  • Example: $Zn(s) + 2 HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$ (Occurs).
  • Example: $Cu(s) + HCl(aq) \rightarrow \text{No Reaction}$.