Ionic and net ionic equations Knowledge-Based Study

Study notes

1. Understanding Aqueous Reactions

   Many chemical reactions take place in aqueous solution, where water serves as the solvent. When soluble ionic compounds dissolve in water, their constituent ions separate from the crystal lattice and move freely, surrounded by water molecules. These substances are called strong electrolytes.

   A traditional chemical equation is known as a molecular equation. In a molecular equation, all reactants and products are written as neutral, intact chemical formulas regardless of whether they exist as free ions in solution.

   For example, mixing aqueous solutions of silver nitrate and sodium chloride produces a precipitate of insoluble silver chloride along with soluble sodium nitrate:

   $$\text{AgNO}_3(aq) + \text{NaCl}(aq) \rightarrow \text{AgCl}(s) + \text{NaNO}_3(aq)$$

   While this molecular representation helps determine stoichiometry and mole relationships, it does not show the actual physical state of the dissolved ions.

   Writing the Complete Ionic Equation

   A complete ionic equation (or total ionic equation) reveals the true species present in solution by showing all soluble strong electrolytes dissociated into their respective aqueous ions.

   To write a complete ionic equation correctly, follow these rules regarding physical states:

   • Aqueous strong electrolytes (aq): Dissociate soluble salts, strong acids, and strong bases completely into their cation and anion components. Remember to multiply ion subscripts by stoichiometric coefficients.
   • Solids (s): Precipitates remain as intact neutral formulas because they form an insoluble solid lattice.
   • Liquids (l) and Gases (g): Pure liquids (such as water) and gases (such as carbon dioxide) do not dissociate and remain in molecular form.
   • Weak electrolytes: Weak acids and weak bases dissociate only slightly in solution and are kept as intact molecular formulas.

   Applying these rules to our previous example:

   $$\text{Ag}^+(aq) + \text{NO}_3^-(aq) + \text{Na}^+(aq) + \text{Cl}^-(aq) \rightarrow \text{AgCl}(s) + \text{Na}^+(aq) + \text{NO}_3^-(aq)$$

   Identifying Spectator Ions

   Notice that in the complete ionic equation, certain ions appear completely identical on both the reactant and product sides. In our example, $\text{Na}^+(aq)$ and $\text{NO}_3^-(aq)$ are present as free ions before and after the reaction takes place.

   Ions that do not participate directly in the chemical change are called spectator ions. They remain in solution throughout the process without undergoing any chemical transformation.

   Deriving the Net Ionic Equation

   A net ionic equation includes only the particles that directly take part in the chemical reaction, leaving out the spectator ions. To determine the net ionic equation:

   1. Write and balance the molecular equation.
   2. Write the complete ionic equation by dissociating all strong electrolytes in aqueous solution.
   3. Cancel identical spectator ions appearing on both sides of the equation.
   4. Verify that both mass (atoms of each element) and net electric charge are balanced.

   Canceling $\text{Na}^+(aq)$ and $\text{NO}_3^-(aq)$ gives the net ionic equation:

   $$\text{Ag}^+(aq) + \text{Cl}^-(aq) \rightarrow \text{AgCl}(s)$$

   Example 1: Precipitation Reaction

   Consider the reaction between aqueous barium chloride and aqueous sodium sulfate:

   Molecular equation:
   $$\text{BaCl}_2(aq) + \text{Na}_2\text{SO}_4(aq) \rightarrow \text{BaSO}_4(s) + 2\,\text{NaCl}(aq)$$

   Complete ionic equation:
   $$\text{Ba}^{2+}(aq) + 2\,\text{Cl}^-(aq) + 2\,\text{Na}^+(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s) + 2\,\text{Na}^+(aq) + 2\,\text{Cl}^-(aq)$$

   Spectator ions: $\text{Na}^+(aq)$ and $\text{Cl}^-(aq)$

   Net ionic equation:
   $$\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$$

   Example 2: Strong Acid-Strong Base Neutralization

   When hydrochloric acid reacts with sodium hydroxide:

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

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

   Net ionic equation:
   $$\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$$

   Every neutralization reaction between a strong monoprotic acid and a soluble strong base reduces to this fundamental net ionic equation.