Net Ionic Reactions
High-Yield Summary
- Net ionic equations split aqueous compounds into ions (keeping solids together) and drop spectator ions — those unchanged on both sides.
- Combination (H₂+F₂→2HF) and decomposition (ammonium dichromate) reactions can be redox, analyzed as one oxidation + one reduction half-reaction.
- Combustion reactions (e.g. CH₄+2O₂→CO₂+2H₂O) are redox where the net ionic equation equals the overall equation — no aqueous/spectator species.
- Double-displacement (metathesis) reactions are usually NOT redox — ions keep their oxidation states; only the ions that actually combine appear (Ag⁺+Cl⁻→AgCl(s)).
- Disproportionation: the same element is simultaneously oxidized and reduced in one reaction — e.g. catalase: 2H₂O₂→2H₂O+O₂.
- Redox titrations use electron transfer instead of protons; iodimetric titrations use iodine + starch indicator, potentiometric titrations track voltage with no indicator.
Key Terms
- Net ionic equation
- Equation with aqueous compounds split into constituent ions, solids kept whole, and spectator ions omitted.
- Spectator ion
- An ion that appears unchanged on both sides of a reaction and doesn't participate in the actual chemistry.
- Disproportionation (dismutation)
- A redox reaction where a single element is simultaneously oxidized and reduced, producing two different products.
- Iodimetric titration
- Redox titration using iodine as titrant; detects endpoint via a blue-black starch–iodine complex.
- Potentiometric titration
- Redox titration using a voltmeter (not a chemical indicator) to track voltage change and find the equivalence point.
Reaction Types Through a Redox Lens
| Reaction type | Redox behavior |
|---|---|
| Combination (H₂+F₂→2H⁺+2F⁻) | Redox — H oxidized (0→+1), F reduced (0→-1); no spectator ions |
| Decomposition ((NH₄)₂Cr₂O₇) | Redox — N oxidized (-3→0), Cr reduced (+6→+3); no spectator ions |
| Combustion (CH₄+2O₂→CO₂+2H₂O) | Redox — C oxidized (-4→+4), O reduced (0→-2); net ionic = overall equation |
| Double-displacement (AgNO₃+HCl) | Usually NOT redox — ions keep oxidation states; net: Ag⁺+Cl⁻→AgCl(s) |
Disproportionation Example: Catalase Breaking Down H₂O₂
- 1Reaction: 2H₂O₂(aq) → 2H₂O(l) + O₂(g)
- 2In H₂O₂, each oxygen atom starts at oxidation state -1.
- 3One oxygen is oxidized: -1 → 0, forming O₂ (loses electrons).
- 4The other oxygen is reduced: -1 → -2, forming H₂O (gains electrons).
- 5Same element (oxygen), both directions at once — the hallmark of disproportionation.
Must-Know Points
- Disproportionation shows up frequently in biological systems — catalase protecting cells from oxidative damage via H₂O₂ breakdown is the classic MCAT example.
- Double-displacement reactions typically form a precipitate, gas, or weak electrolyte — not electron transfer — which is why they're usually not redox.
Common MCAT Trap
- Don't assume every reaction type is redox by default — double-displacement (metathesis) reactions are the main exception since oxidation states don't change.
- In disproportionation, the SAME element ends up in two different products at two different oxidation states — don't mistake it for a normal two-species redox reaction.
Quick Recall
What happens to spectator ions in a net ionic equation?
Is a double-displacement reaction usually a redox reaction?
What indicator is used in an iodimetric titration, and what color change signals the endpoint?