Half Equations & Ionic Equations

Writing Net Ionic Equations

• Step 1: Write the full balanced molecular equation including state symbols for all reactants and products.
• Step 2: Dissociate all aqueous (aq) ionic compounds into their constituent ions; do not dissociate solids, liquids, or gases.
• Step 3: Identify and cancel spectator ions that appear identically on both sides of the equation.
• Step 4: Rewrite the remaining species to form the net ionic equation, ensuring both atoms and charges are balanced.

Constructing Half Equations

• Step 1: Identify the species undergoing change (e.g., a metal atom becoming a metal ion).
• Step 2: Balance the atoms of the element being oxidized or reduced on both sides.
• Step 3: Calculate the total charge on each side and add electrons ($e^-$) to the more positive side to equalize the charge.
• Step 4: Verify that oxidation shows electrons as products (right side) and reduction shows electrons as reactants (left side).

• The primary difference between a molecular equation and an ionic equation is the level of detail regarding the state of substances in solution.

• While both require balancing, a half equation is an incomplete chemical event that must be paired with another half equation to represent a real-world reaction.

• Check State Symbols: Examiners often award marks for correct state symbols; remember that only $(aq)$ species are split into ions, while $(s)$, $(l)$, and $(g)$ are never split.

• The Charge Balance Rule: In a net ionic equation, the total charge on the left does not have to be zero, but it must equal the total charge on the right.

• Electron Placement: Use the mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain) to place electrons; if a species is oxidized, electrons must be on the right side of the arrow.

• Simplifying Ratios: Always ensure the coefficients in your final ionic equation are in the simplest whole-number ratio possible.

• Splitting Solids: A common error is splitting precipitates like $AgCl(s)$ into ions; these must remain as a single formula unit because they are not dissociated in the solvent.

• Ignoring Subscripts: When splitting $MgCl_2(aq)$, students often write $Cl_2^-$ instead of $2Cl^-$; remember that subscripts in a formula become coefficients when the substance dissociates.

• Unbalanced Charges: Students frequently balance the atoms but forget to check if the total charge is the same on both sides, which is a critical requirement for a valid equation.

• Spectator Confusion: Assuming all ions are spectators is incorrect; only those that remain in the $(aq)$ state and do not change oxidation number are spectators.