Writing Balanced Ionic Equations

Step 1: Write the Balanced Molecular Equation: Ensure all reactants and products are correctly identified and the equation is balanced using stoichiometric coefficients.

Step 2: Determine Solubility: Use solubility rules to identify which compounds are soluble $(aq)$ and which are insoluble $(s)$, liquids $(l)$, or gases $(g)$.

Step 3: Dissociate Strong Electrolytes: Rewrite the equation by breaking all $(aq)$ strong electrolytes into their constituent ions. Multiply the ion's subscript by the stoichiometric coefficient to get the total count.

Step 4: Identify and Cancel Spectators: Cross out ions that appear exactly the same (same charge, same state, same quantity) on both sides of the equation.

Step 5: Finalize the Net Ionic Equation: Write the remaining species and perform a final check for mass and charge balance.

The Charge Check: Always verify that the total charge on the left equals the total charge on the right. If the left side is $+2$ and the right side is $0$, the equation is incorrect.

Subscript vs. Coefficient: When dissociating $CaCl_2$, the subscript '2' becomes a coefficient: $Ca^{2+} + 2Cl^-$. Do not write $Cl_2^{2-}$.

Identify the 'Driver': Look for the formation of a solid, a gas, or a weak electrolyte (like water). If all reactants and products are soluble ions, 'No Reaction' occurs.

State Symbol Precision: Marks are often lost for omitting $(aq)$, $(s)$, $(l)$, or $(g)$. These symbols are the primary indicators of which species should be dissociated.

Dissociating Solids: A common error is breaking apart a precipitate into ions. Solids, liquids, and gases must remain as whole molecules/units.

Ignoring Polyatomic Ions: Students often break polyatomic ions into individual elements. For example, $SO_4^{2-}$ must stay as a single unit, not $S^{6+} + 4O^{2-}$.

Incorrect Ion Charges: Forgetting the specific charge of common ions (e.g., writing $Ag^+$ as $Ag^{2+}$) will lead to an unbalanced charge and an incorrect final equation.