USNCO Reaction Writing
163 questions from 1999–2025 with answers, justifications, and cross-references.
Writing reactions requires reasoning from first principles — driving forces, oxidation states, and acid-base hierarchies.
Official USNCO Instructions
"Write a balanced, net ionic equation for the reaction that occurs. Represent substances in solution as ions if the substances are extensively ionized. Omit formulas for any ions or molecules that are unchanged by the reaction. You need not balance the equations."
First Principles Reference
| # | Principle | Summary |
|---|---|---|
| FP1 | Driving Forces | Reactions proceed toward precipitates, gases, weak electrolytes (H₂O), stable complexes, or more stable oxidation states. |
| FP2 | Oxidation State Tracking | Assign oxidation states → identify what is oxidized/reduced → electron balance determines stoichiometry. |
| FP3 | Acid-Base Hierarchy | Stronger acid displaces weaker acid; compare Kₐ values to predict which proton transfers are favorable. |
| FP4 | Lewis Structure / Electron Availability | Lone pairs donate to empty orbitals; nucleophiles attack electrophiles. |
| FP5 | Periodic Trends | Charge density (Z²/r), electronegativity, metallic character, E° predict chemical behavior from periodic position. |
| FP6 | Conservation Laws | Mass, charge, and electrons are conserved; these constrain and verify balanced equations. |
Common Oxidizing & Reducing Agents
Redox is the largest category (~50 questions). Knowing what each agent produces — including its diagnostic color change — is the key to writing these equations.
Oxidizing Agents (get reduced)
| Agent | Condition | Product | e⁻ change | Key detail |
|---|---|---|---|---|
| MnO₄⁻ | acidic | Mn²⁺ | +7→+2 (5e⁻) | Purple → colorless |
| MnO₄⁻ | neutral/basic | MnO₂ | +7→+4 (3e⁻) | Purple → brown ppt |
| Cr₂O₇²⁻ | acidic | Cr³⁺ | +6→+3 (6e⁻/ion) | Orange → green. Oxidizes alcohols |
| NO₃⁻ (HNO₃) | dilute | NO | +5→+2 (3e⁻) | Colorless gas, browns in air |
| NO₃⁻ (HNO₃) | concentrated | NO₂ | +5→+4 (1e⁻) | Brown gas |
| H₂SO₄ | conc., hot | SO₂ | +6→+4 (2e⁻) | Dissolves Cu. Molecular equation |
| H₂O₂ | as oxidizer | H₂O | −1→−2 (2e⁻) | Oxidizes I⁻ → I₂ |
| MnO₂ | + conc. HCl | Mn²⁺ | +4→+2 (2e⁻) | Lab prep of Cl₂ (3× on USNCO) |
Reducing Agents (get oxidized)
| Agent | Product | e⁻ change | Key detail |
|---|---|---|---|
| Fe²⁺ | Fe³⁺ | +2→+3 (1e⁻) | Classic titration with MnO₄⁻ or Cr₂O₇²⁻ |
| Sn²⁺ | Sn⁴⁺ | +2→+4 (2e⁻) | Paired with Cr₂O₇²⁻ |
| I⁻ | I₂ | −1→0 (1e⁻) | Oxidized by H₂O₂, Cu²⁺, IO₃⁻ |
| C₂O₄²⁻ | CO₂ | +3→+4 (2e⁻/ion) | Paired with KMnO₄. Gas escapes |
| SO₂ / SO₃²⁻ | SO₄²⁻ | +4→+6 (2e⁻) | Mild reductant |
| S₂O₃²⁻ | S₄O₆²⁻ | +2→+2.5 (1e⁻/ion) | Titrant for I₂ (iodometric titration) |
| Cl⁻ (conc.) | Cl₂ | −1→0 (1e⁻) | Needs conc. HCl + strong oxidizer |
| RCH₂OH (1° alcohol) | RCHO / RCOOH | 2e⁻ or 4e⁻ loss | PCC stops at aldehyde; Cr₂O₇²⁻ or KMnO₄ gives full oxidation to acid |
| R₂CHOH (2° alcohol) | R₂C=O | 2e⁻ loss | Stops at ketone — no further oxidation |
Special Redox Patterns
Comproportionation
Two different oxidation states of the same element converge to one intermediate state.
e.g. ClO⁻(+1) + Cl⁻(−1) → Cl₂(0)
6 questions on USNCO
Disproportionation
One oxidation state splits into two — one higher, one lower.
e.g. Cl₂(0) → ClO⁻(+1) + Cl⁻(−1)
Reverse of comproportionation
Displacement
A more reactive element displaces a less reactive one. Check the activity series / halogen order.
e.g. Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
8 questions on USNCO
Solubility Rules for Precipitation
Precipitation accounts for ~14 questions. Knowing what's insoluble is essential.
Generally SOLUBLE (dissolve)
- All Group 1 (Na⁺, K⁺, Li⁺) and NH₄⁺ salts — mostly soluble, mostly spectators
- All NO₃⁻ (nitrate) salts — no exceptions
- All CH₃COO⁻ (acetate) salts — AgCH₃COO is sparingly soluble
- All Cl⁻, Br⁻, I⁻ (halide) salts — EXCEPT AgX, PbX₂, Hg₂X₂
- All F⁻ (fluoride) salts — EXCEPT CaF₂, BaF₂, MgF₂, PbF₂
- All SO₄²⁻ (sulfate) salts — EXCEPT BaSO₄, PbSO₄, SrSO₄
Generally INSOLUBLE (precipitate)
- All OH⁻ (hydroxide) salts — EXCEPT Group 1, Ba(OH)₂, Sr(OH)₂; Ca(OH)₂ slightly
- All CO₃²⁻ (carbonate) salts — EXCEPT Group 1, NH₄⁺
- All PO₄³⁻ (phosphate) salts — EXCEPT Group 1, NH₄⁺
- All S²⁻ (sulfide) salts — EXCEPT Group 1, Group 2, NH₄⁺
- All CrO₄²⁻ (chromate) salts — EXCEPT Group 1, NH₄⁺; BaCrO₄, PbCrO₄, Ag₂CrO₄ insoluble
Most-Tested Precipitates on USNCO
| Precipitate | Ksp | Appearances | Notes |
|---|---|---|---|
| BaSO₄ | 1.1×10⁻¹⁰ | 5× | Most common. Often in double ppt with Mg(OH)₂ |
| Mg(OH)₂ | 5.6×10⁻¹² | 4× | Paired with BaSO₄ in double ppt questions |
| AgCl / AgBr | 1.8×10⁻¹⁰ / 5.4×10⁻¹³ | 6× | Dissolved by NH₃ or S₂O₃²⁻ (complexation) |
| PbSO₄ | 2.5×10⁻⁸ | 2× | One of the rare insoluble sulfates |
| PbCrO₄ | 2×10⁻¹⁴ | 1× | Bright yellow "chrome yellow" pigment |
| Ag₂CrO₄ | 1.1×10⁻¹² | 1× | Brick-red. Mohr titration endpoint indicator |
| Hg₂Cl₂ | 1.4×10⁻¹⁸ | 1× | White. Hg(I) always as Hg₂²⁺ dimer |
| CaF₂ | 3.5×10⁻¹¹ | 1× | Precipitate-to-precipitate (CaCO₃ → CaF₂) |
| NiS | ~10⁻²⁰ | 1× | Ppts even in acid (Ksp overwhelms H⁺) |
| CuI | 1.3×10⁻¹² | 1× | Precipitation drives Cu²⁺ + I⁻ redox |
Reaction Type Guide
Proton transfer reactions including strong/weak acid-base, neutralization, amphoteric reactions, acid anhydride hydrolysis, and non-redox thermal decompositions (carbonates, sulfites).
Formation of insoluble products from mixing ionic solutions, including double precipitation and precipitate-to-precipitate conversions.
Electron transfer reactions including displacement, strong oxidizers (KMnO₄, K₂Cr₂O₇, HNO₃), disproportionation, comproportionation, thermal decomposition, and electrolysis.
Lewis acid-base reactions forming coordination complexes, including ammine, cyanide, thiosulfate, hydroxo, halide, and thiocyanate complexes.
Carbon chemistry including electrophilic addition, EAS (nitration/halogenation), elimination (E1/E2), substitution (SN1/SN2), saponification, and polymerization.
Nuclear transformations including alpha decay, beta decay, positron emission, electron capture, fission, and nuclear bombardment.
Hydrolysis reactions of oxides, ionic compounds, covalent halides, and peroxides in water, including thermal decompositions of salts.
How to Approach Unknown Reactions
Read & Decode
Names → formulas. This is where many students get stuck. If you can't write the formula, you can't write the equation.
Watch for keywords that change the product:
"excess" — Different product (HCO₃⁻ not CO₃²⁻, complex not ppt)"concentrated" — HNO₃→NO₂, HCl as reductant, H₂SO₄ as oxidizer"dilute" — HNO₃→NO, acid as proton source only"heated strongly" — Thermal decomposition, complete breakdownClassify & Predict Products
First, two easy checks:
- Nuclear? Isotope notation, decay particles (α, β, positron, neutron). Use mass number balance and atomic number balance to find the unknown nuclide.
- Organic? C–H bonds, functional groups. Identify the mechanism (addition, elimination, EAS, SN1/SN2, saponification). Don't miss the small molecule byproduct (H₂O, HCl, ROH).
For inorganic reactions — the key question: Redox or Non-redox?
Look for common oxidizing agents (MnO₄⁻, Cr₂O₇²⁻, HNO₃, conc. H₂SO₄) or reducing agents (metals, I⁻, Fe²⁺, SO₃²⁻). If oxidation states change → redox.
If REDOX:
- Identify the oxidizer and reducer — use the reference table above to know what each produces
- Determine products based on common oxidation states and the medium: Acidic → use H⁺ and H₂O to balance. Basic/neutral → MnO₄⁻ gives MnO₂ not Mn²⁺
- Electron balance — e⁻ lost = e⁻ gained → determines stoichiometric coefficients
- Add H₂O to whichever side needs oxygen atoms balanced
If NON-REDOX — further classify by driving force:
- Precipitation — Two ionic solutions mixed → check solubility rules → insoluble product forms. Watch for double precipitation.
- Acid-Base — Proton transfer; stronger acid displaces weaker. Compare Kₐ values for polyprotic species. Check for double driving forces.
- Ext. Hydrolysis — Solid oxide, halide, or ionic compound + water → oxyacid/base, nitride/carbide → NH₃/C₂H₂, covalent halide → HX + oxyacid.
- Complexation — Ligand coordinates to metal ion. Common ligands (NH₃, CN⁻, OH⁻, SCN⁻, S₂O₃²⁻, Cl⁻, F⁻, C₂O₄²⁻). "Excess" is the clue. HSAB theory: hard acids (Fe³⁺, Al³⁺) prefer hard bases (F⁻, OH⁻); soft acids (Ag⁺, Cu⁺) prefer soft bases (CN⁻, S²⁻, I⁻).
Some reactions involve multiple types (precipitation + acid-base, redox driven by precipitation as in Cu²⁺ + I⁻ → CuI + I₂). Always check for a second driving force.
Write Net Ionic
- As ions: strong acids, strong bases, soluble salts
- As molecules: weak acids/bases, H₂O, gases, organic compounds
- As formulas: solids (precipitates, undissolved reactants), pure liquids
- Omit: spectator ions that appear unchanged on both sides
Balance & Check
- Atoms: every element balanced on both sides
- Charge: total charge equal on both sides
- Electrons: (redox only) e⁻ lost = e⁻ gained
- Driving force: at least one present — if not, reconsider
Note: USNCO Part II does not require balancing or phase labels — but including them deepens your understanding and helps catch errors.
Common Mistakes
Concrete wrong → right examples drawn from actual USNCO questions.
Ionic vs. Molecular
Writing weak electrolytes as ions
HF → H⁺ + F⁻ · CH₃COOH → CH₃COO⁻ + H⁺ · NH₃ → NH₄⁺ + OH⁻HF(aq) stays molecular · CH₃COOH(aq) stays molecular · NH₃(aq) stays molecularOnly the 6 strong acids (HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄ 1st proton), strong bases (Group 1/2 hydroxides), and soluble salts are written as ions. Everything else stays molecular.
Dissociating an undissolved solid
Na₂SO₃(s) written as 2Na⁺ + SO₃²⁻Na₂SO₃(s) — the solid has not dissolved, so it must be written as the formula unitIf a reactant is described as 'solid,' it hasn't dissolved yet. Write it as the complete formula with (s). It only becomes ions after it dissolves.
Reading the Problem
Ignoring 'excess' — it changes the product
Excess CO₂ + Ca(OH)₂ → CaCO₃ (treating it like limited CO₂)Excess CO₂ + Ca(OH)₂ → Ca(HCO₃)₂ — the excess CO₂ converts CO₃²⁻ to HCO₃⁻'Excess' is never a throwaway word. Excess CO₂ → bicarbonate not carbonate. Excess NH₃ → ammine complex not hydroxide precipitate. Excess OH⁻ → aluminate/zincate complex not hydroxide precipitate. Always ask: what does the excess reagent do to the initial product?
Not recognizing the compound from its name
Stuck on 'silica' or 'nitrosyl fluoride' or 'propyl benzoate'silica = SiO₂ · nitrosyl fluoride = NOF · propyl benzoate = C₆H₅COOC₃H₇ · calcination = heating to decomposeDecoding the name is often the hardest step. Key vocabulary: -ide (binary), -ite/-ate (oxyanions), -ous/-ic (acid strength), propyl/butyl (carbon count), -oate/-yl (ester = [alcohol] [acid]-ate). If you can't write the formula, you can't write the equation.
Redox Errors
Forgetting H⁺ and H₂O in acidic redox reactions
MnO₄⁻ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ (unbalanced — where do the oxygens go?)MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂OIn acidic solution, the oxygens in the oxidizer (MnO₄⁻, Cr₂O₇²⁻, NO₃⁻) must be accounted for: they combine with H⁺ to form H₂O. Balance O with H₂O, then balance H with H⁺, then verify charge.
Writing only half of a redox reaction
2Al + 2OH⁻ + 6H₂O → 2[Al(OH)₄]⁻ (only shows Al oxidized — where is the reduction? H₂ is missing!)2Al + 2OH⁻ + 6H₂O → 2[Al(OH)₄]⁻ + 3H₂ (Al oxidized 0→+3, H reduced +1→0 in H₂)Electrons cannot appear in a net equation. If something is oxidized, something else must be reduced. In Al + NaOH, Al is oxidized but students forget that H₂O provides the H⁺ that gets reduced to H₂ gas. Use electron balance: e⁻ lost = e⁻ gained.
Wrong product for the oxidizing agent
MnO₄⁻ in neutral solution → Mn²⁺ (that's the acidic product)MnO₄⁻ in neutral/basic → MnO₂(s) · MnO₄⁻ in acid → Mn²⁺The reduction product depends on conditions. MnO₄⁻: acid → Mn²⁺ (5e⁻), neutral → MnO₂ (3e⁻). HNO₃: dilute → NO (3e⁻), conc. → NO₂ (1e⁻). See the oxidizing agents reference table above.
Acid-Base Errors
Missing one net-ionic equation out of two
Pb²⁺ + 2Br⁻ → PbBr₂(s) only (wrote precipitation but missed the weak acid formation)Pb²⁺ + 2CH₃COO⁻ + 2H⁺ + 2Br⁻ → PbBr₂(s) + 2CH₃COOH — both driving forces includedMany USNCO reactions have TWO simultaneous driving forces: precipitation + weak electrolyte formation, or acid-base + gas evolution. If you write only one, you'll have 'leftover' spectator ions that actually participate. Check: after writing your equation, are there ions on both sides that could react with each other? If CH₃COO⁻ and H⁺ are both 'spectators,' they aren't — they form CH₃COOH.
Wrong degree of protonation with polyprotic species
CH₃COOH + PO₄³⁻ → CH₃COO⁻ + H₃PO₄ (acetic acid can't protonate all the way)CH₃COOH + PO₄³⁻ → only to HPO₄²⁻ or H₂PO₄⁻ — compare Kₐ values at each stepFor polyprotic species, check each protonation step: K = Kₐ(acid)/Kₐ(conjugate). If K ≫ 1, the step proceeds. If K ≪ 1, it stops. CH₃COOH (Kₐ=1.8×10⁻⁵) can protonate PO₄³⁻ → HPO₄²⁻ (K=4×10⁷ ✓) and HPO₄²⁻ → H₂PO₄⁻ (K=286 ✓) but NOT H₂PO₄⁻ → H₃PO₄ (K=0.003 ✗).
Organic Errors
Forgetting small molecule byproducts in organic reactions
CH₃CH(OH)CH₃ → CH₃CH=CH₂ (where did the OH go?)CH₃CH(OH)CH₃ → CH₃CH=CH₂ + H₂O — water is the byproduct of eliminationOrganic reactions almost always produce a small molecule byproduct. Elimination → H₂O. EAS (nitration) → H₂O. Saponification → ROH (alcohol). Ester hydrolysis → the alcohol fragment. Addition of HX or Br₂ → no byproduct (everything adds). Polymerization (condensation) → HCl or H₂O per monomer unit. If atoms are missing from your product, you forgot the byproduct.
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