Naming & Reactions

Two skills that first-semester chemistry keeps asking for and rarely teaches in one place: turning a formula into a name (and back), and looking at a set of reactants and predicting what they make. Both are procedures. This page lays out the procedures as decision steps with worked examples, and ends with a predict-the-products drill and answers.

Use it with the Quick Reference (the ion, solubility, and activity-series tables) and the Problem-Solving Playbook.

Part A — Naming compounds

First, decide what kind of compound it is

Look at the formula (or the name) and sort it into one bucket. The naming rules are different for each, so this step is not optional.

• Ionic — a metal with a nonmetal (NaCl, CaBr₂), or anything containing a polyatomic ion (Na₂SO₄, NH₄Cl). The name has no prefixes.
• Molecular (covalent) — two nonmetals only (CO₂, N₂O₄, PCl₅). The name uses prefixes.
• Acid — a formula written with H first and dissolved in water (HCl(aq), H₂SO₄(aq)). Acids have their own naming rules.
• Hydrate — an ionic formula with ·nH₂O attached (CuSO₄·5H₂O).

A metal-plus-nonmetal is always ionic even if the electronegativity difference is modest; two nonmetals are always molecular. That single distinction — is there a metal? — decides whether you use prefixes.

Ionic compounds

Rule: (name of the cation) (name of the anion), with the anion ending changed to -ide for a single element. No prefixes; the subscripts are implied by the charges.

1. Cation — the metal (or NH₄⁺, ammonium). Keep its element name: sodium, calcium, aluminum.
2. If the metal is a transition metal or a p-block metal with more than one common charge (iron, copper, tin, lead, chromium, and most others past group 2), you must show the charge with a Roman numeral: iron(II), iron(III), copper(I), copper(II), lead(II). Work the charge out from the anion. The always-one-charge metals that need no numeral: group 1 (+1), group 2 (+2), aluminum (+3), zinc (+2), silver (+1), cadmium (+2).
3. Anion — a single nonmetal becomes -ide: chloride, oxide, sulfide, nitride, phosphide. A polyatomic ion keeps its own name: sulfate, nitrate, carbonate, hydroxide, phosphate.

Worked:

• CaCl₂ → calcium is always +2, so no numeral → calcium chloride.
• FeCl₃ → three Cl⁻ means the iron is +3 → iron(III) chloride.
• Fe₂O₃ → three O²⁻ (charge −6) balanced by two Fe (each +3) → iron(III) oxide.
• Cu₂S → one S²⁻ balanced by two Cu (each +1) → copper(I) sulfide.
• (NH₄)₃PO₄ → ammonium and phosphate → ammonium phosphate.
• Mg(OH)₂ → magnesium hydroxide (not "dihydroxide" — no prefixes in ionic names).

Going backward (name → formula): write each ion with its charge, then use the crossover: the numerical charge of one ion becomes the subscript of the other. Reduce to lowest whole numbers. "Aluminum sulfide" → Al³⁺ and S²⁻ → cross → Al₂S₃. "Calcium phosphate" → Ca²⁺ and PO₄³⁻ → cross → Ca₃(PO₄)₂ (parentheses around the polyatomic ion when there is more than one).

The polyatomic ions to memorize

There is no shortcut for the core set; memorize these. Grouped so the patterns help:

The -ate / -ite pattern: for the same central atom, -ate has one more oxygen than -ite, and both keep the same charge (sulfate SO₄²⁻ / sulfite SO₃²⁻; nitrate NO₃⁻ / nitrite NO₂⁻).

The per- / hypo- pattern (chlorine oxyanions, all −1): perchlorate ClO₄⁻ (one more O than chlorate) → chlorate ClO₃⁻ → chlorite ClO₂⁻ → hypochlorite ClO⁻ (one fewer O than chlorite). Bromine and iodine follow the same series.

Adding H⁺: attaching a hydrogen to a −2 ion makes a −1 ion and adds "hydrogen" to the name: CO₃²⁻ → HCO₃⁻ hydrogen carbonate; PO₄³⁻ → H₂PO₄⁻ dihydrogen phosphate.

Molecular (covalent) compounds

Rule: (prefix)(first element) (prefix)(second element)-ide. Prefixes give the exact number of each atom, because — unlike ionic compounds — there is no charge to imply the ratio.

Two conventions: drop "mono-" on the first element only (CO is carbon monoxide, not monocarbon monoxide), and elide a doubled vowel at a prefix–element boundary (CO → monoxide not monooxide; N₂O₄ → tetroxide).

Worked:

• CO₂ → carbon dioxide.
• N₂O → dinitrogen monoxide (informally "nitrous oxide").
• P₄O₁₀ → tetraphosphorus decoxide.
• SF₆ → sulfur hexafluoride.
• Cl₂O₇ → dichlorine heptoxide.

Backward: the prefixes are the subscripts. "Diphosphorus pentoxide" → P₂O₅.

Acids

An acid is named from the anion you would get by removing the H⁺(s).

• Anion ends in -ide (binary acid, no oxygen): hydro-(root)-ic acid. HCl(aq) → chloride → hydrochloric acid. H₂S(aq) → hydrosulfuric acid.
• Anion ends in -ate (oxyacid): (root)-ic acid. HNO₃ → nitrate → nitric acid. H₂SO₄ → sulfate → sulfuric acid. H₃PO₄ → phosphate → phosphoric acid (the root keeps the "phosphor").
• Anion ends in -ite (oxyacid): (root)-ous acid. HNO₂ → nitrite → nitrous acid. H₂SO₃ → sulfite → sulfurous acid.
• The per-/hypo- prefixes carry over: HClO₄ → perchloric acid; HClO → hypochlorous acid.

Mnemonic for the endings: "-ate → -ic, -ite → -ous" — "I ate something icky; the ite was oust." (Say it once and it sticks.)

Hydrates

Name the ionic compound, then a prefix + "hydrate" for the waters: CuSO₄·5H₂O → copper(II) sulfate pentahydrate. Heating drives the water off, leaving anhydrous copper(II) sulfate; the mass lost is used in a formula-of-a-hydrate lab.

Part B — Predicting the products of a reaction

Given the reactants, decide which of these patterns applies. Recognizing the pattern tells you the products; then you balance.

1. Combination (synthesis) — A + B → AB

Two or more simple substances join into one.

• metal + nonmetal → ionic compound. Predict the formula from the ions. 2 Mg + O₂ → 2 MgO. 2 Na + Cl₂ → 2 NaCl.
• metal oxide + water → metal hydroxide (CaO + H₂O → Ca(OH)₂); nonmetal oxide + water → oxyacid (SO₃ + H₂O → H₂SO₄; CO₂ + H₂O → H₂CO₃).
• two nonmetals → molecular compound, often more than one possible; the problem usually names it (N₂ + 3 H₂ → 2 NH₃).

2. Decomposition — AB → A + B

One compound breaks into simpler pieces. Usually driven by heat or electricity.

• metal carbonate → metal oxide + CO₂ (CaCO₃ → CaO + CO₂).
• metal hydroxide → metal oxide + H₂O (Ca(OH)₂ → CaO + H₂O).
• metal chlorate → metal chloride + O₂ (2 KClO₃ → 2 KCl + 3 O₂).
• binary compound → its elements (2 H₂O → 2 H₂ + O₂ by electrolysis).

3. Combustion — fuel + O₂ → oxides (+ heat)

A hydrocarbon or a C/H/O compound burned in excess oxygen gives CO₂ and H₂O, always. Balance C first, then H, then O last.

C₃H₈ + 5 O₂ → 3 CO₂ + 4 H₂O (propane). 2 C₂H₆O + 6 O₂ → 4 CO₂ + 6 H₂O (ethanol).

Limited oxygen gives CO or soot (C); the problem will say so.

4. Single replacement — A + BC → AC + B

One element displaces another from a compound. Whether it happens depends on the activity series (on the Quick Reference): a metal displaces any metal (or hydrogen) below it on the list; it does nothing to one above it.

• metal + acid → salt + H₂ — if the metal is above hydrogen. Zn + 2 HCl → ZnCl₂ + H₂↑. Cu + HCl → no reaction (Cu is below H).
• metal + salt → new salt + metal — if the free metal is above the one in the salt. Fe + CuSO₄ → FeSO₄ + Cu (iron is above copper). Cu + FeSO₄ → no reaction.
• halogen + halide salt — a halogen displaces any halogen below it in group 17. Cl₂ + 2 NaBr → 2 NaCl + Br₂; Br₂ + NaCl → no reaction.

5. Double replacement (metathesis) — AB + CD → AD + CB

Two ionic compounds in solution swap partners. This only goes if one product leaves the solution — as a precipitate, a gas, or water (a weak electrolyte). If all four ions stay dissolved, no reaction.

Precipitation. Swap the ions, then check each product against the solubility rules:

Soluble: almost all nitrates, acetates, group 1 and ammonium salts; most chlorides/bromides/iodides (except Ag⁺, Pb²⁺, Hg₂²⁺); most sulfates (except Ba²⁺, Pb²⁺, Ca²⁺ slightly, Sr²⁺). Insoluble: most carbonates, phosphates, sulfides, hydroxides (except group 1 and ammonium; Ba(OH)₂ and Ca(OH)₂ are somewhat soluble).

AgNO₃(aq) + NaCl(aq) → AgCl(s)↓ + NaNO₃(aq) — AgCl is insoluble, so the reaction proceeds. NaNO₃(aq) + KCl(aq) → no reaction — both possible products are soluble.

Net ionic equation, step by step:

1. Write the balanced molecular equation with (aq)/(s)/(g)/(l). Pb(NO₃)₂(aq) + 2 KI(aq) → PbI₂(s) + 2 KNO₃(aq)
2. Write the total ionic equation: split every strong electrolyte (soluble ionic compounds, strong acids/bases) into ions; leave solids, gases, water, and weak electrolytes intact. Pb²⁺ + 2 NO₃⁻ + 2 K⁺ + 2 I⁻ → PbI₂(s) + 2 K⁺ + 2 NO₃⁻
3. Cancel the spectator ions (identical on both sides): NO₃⁻ and K⁺.
4. What remains is the net ionic equation: Pb²⁺(aq) + 2 I⁻(aq) → PbI₂(s) Check that atoms and charge balance (left: +2 −2 = 0; right: 0).

Gas-forming. Some swaps produce an unstable product that decomposes to a gas: carbonates or bicarbonates with acid give H₂CO₃, which becomes H₂O + CO₂↑; sulfites give SO₂; sulfides give H₂S; ammonium salts with strong base give NH₃. CaCO₃(s) + 2 HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)↑

6. Acid–base neutralization — acid + base → salt + water

A special, always-goes double replacement. The H⁺ from the acid and the OH⁻ from the base make water; the leftover ions make the salt.

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) H₂SO₄(aq) + 2 KOH(aq) → K₂SO₄(aq) + 2 H₂O(l) Net ionic for any strong acid + strong base: H⁺(aq) + OH⁻(aq) → H₂O(l).

7. Oxidation–reduction (redox) — electrons move

Any reaction where oxidation numbers change. All combustion and all single-replacement reactions are redox; most combination and decomposition reactions are; double replacement (precipitation, neutralization) generally is not — the ions keep their charges.

Assign oxidation numbers with these rules, in order of priority:

1. A free element is 0 (Na, O₂, P₄).
2. A monatomic ion equals its charge (Na⁺ is +1, S²⁻ is −2).
3. In compounds: F is −1; group 1 is +1, group 2 is +2; H is +1 (−1 with a metal, in hydrides); O is −2 (−1 in peroxides like H₂O₂).
4. The oxidation numbers in a neutral compound sum to 0; in a polyatomic ion they sum to the ion charge. Use this to find the unknown one.

Worked: in KMnO₄, K is +1 and each O is −2 (total −8); the sum is 0, so Mn = +7. In Cr₂O₇²⁻, seven O give −14; the sum is −2, so 2 Cr = +12 and each Cr = +6.

Then: the species whose oxidation number goes up is oxidized (it lost electrons, it is the reducing agent); the one that goes down is reduced (gained electrons, the oxidizing agent). In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc goes 0 → +2 (oxidized) and copper goes +2 → 0 (reduced).

Predict-the-products drill

For each, name the reaction type, give the products with states, and balance. Answers below.

1. Al(s) + O₂(g) →
2. C₄H₁₀(g) + O₂(g) → (complete combustion)
3. Ba(NO₃)₂(aq) + Na₂SO₄(aq) →
4. Zn(s) + AgNO₃(aq) →
5. NaHCO₃(aq) + HC₂H₃O₂(aq) → (baking soda + vinegar)
6. NiCl₂(aq) + KOH(aq) →
7. Ca(s) + H₂O(l) →
8. KI(aq) + Pb(NO₃)₂(aq) → — also give the net ionic equation.
9. H₃PO₄(aq) + Ca(OH)₂(aq) →
10. Cu(s) + HCl(aq) →

Answers

1. Combination (redox). 4 Al(s) + 3 O₂(g) → 2 Al₂O₃(s). Al³⁺ and O²⁻ cross to Al₂O₃.
2. Combustion. 2 C₄H₁₀(g) + 13 O₂(g) → 8 CO₂(g) + 10 H₂O(g).
3. Double replacement — precipitation. Products BaSO₄ and NaNO₃. Sulfate with Ba²⁺ is insoluble → Ba(NO₃)₂(aq) + Na₂SO₄(aq) → BaSO₄(s)↓ + 2 NaNO₃(aq).
4. Single replacement (redox). Zn is above Ag, so it displaces silver: Zn(s) + 2 AgNO₃(aq) → Zn(NO₃)₂(aq) + 2 Ag(s).
5. Double replacement — gas-forming. Carbonate + acid → H₂CO₃ → H₂O + CO₂. NaHCO₃(aq) + HC₂H₃O₂(aq) → NaC₂H₃O₂(aq) + H₂O(l) + CO₂(g)↑.
6. Double replacement — precipitation. Ni(OH)₂ is an insoluble hydroxide: NiCl₂(aq) + 2 KOH(aq) → Ni(OH)₂(s)↓ + 2 KCl(aq).
7. Single replacement (redox). An active metal + water → metal hydroxide + H₂: Ca(s) + 2 H₂O(l) → Ca(OH)₂(aq) + H₂(g)↑.
8. Double replacement — precipitation. 2 KI(aq) + Pb(NO₃)₂(aq) → PbI₂(s)↓ + 2 KNO₃(aq). Net ionic: Pb²⁺(aq) + 2 I⁻(aq) → PbI₂(s) (K⁺ and NO₃⁻ are spectators).
9. Neutralization. 2 H₃PO₄(aq) + 3 Ca(OH)₂(aq) → Ca₃(PO₄)₂(s)↓ + 6 H₂O(l) — calcium phosphate is insoluble, so this one both neutralizes and precipitates.
10. No reaction. Copper is below hydrogen on the activity series, so it does not displace H⁺ from a non-oxidizing acid like HCl.

Where this connects

Naming is the gateway to everything symbolic: you cannot write a formula unit for a molar mass calculation or a balanced equation if you cannot get from "calcium phosphate" to Ca₃(PO₄)₂. Reaction prediction is what makes a stoichiometry problem solvable — the balanced equation that anchors the mole ratio has to come from somewhere. If either skill is shaky, it will not feel like a naming problem or a prediction problem later; it will feel like "I can't do stoichiometry." Fix it here. See How the Ideas Connect for the full dependency map.