Reaction Type Predictor
Enter a chemical equation and identify combustion, synthesis, decomposition, replacement and acid-base patterns.
Reaction patterns checked
| Type | Pattern | Example |
|---|---|---|
| Combustion | Fuel + O2 -> CO2 + H2O | CH4 + O2 -> CO2 + H2O |
| Synthesis | A + B -> AB | 2Na + Cl2 -> 2NaCl |
| Decomposition | AB -> A + B | CaCO3 -> CaO + CO2 |
| Single replacement | A + BC -> AC + B | Zn + CuSO4 -> ZnSO4 + Cu |
| Double replacement | AB + CD -> AD + CB | AgNO3 + NaCl -> AgCl + NaNO3 |
| Neutralization | Acid + base -> salt + water | HCl + NaOH -> NaCl + H2O |
Rules used
Worked examples
CH4 + O2 -> CO2 + H2O: a hydrocarbon reacts with oxygen and forms carbon dioxide and water, so it is combustion.
AgNO3 + NaCl -> AgCl + NaNO3: two ionic compounds exchange ions, so it is double replacement. Solubility rules show AgCl is a precipitate.
Zn + CuSO4 -> ZnSO4 + Cu: one metal replaces another metal ion, so it is single replacement. The activity series helps decide whether it occurs.
How to read the pattern diagram
The visual separates the left and right sides of the equation and labels the likely movement: joining, splitting, swapping, replacing or burning in oxygen. If the predictor also balances the equation, use that balanced form before doing mole ratios or limiting reagent work.
Where this predictor is useful
- Sorting homework equations before writing products or balancing.
- Preparing for reaction classification quizzes.
- Checking whether an equation should use solubility rules or the activity series next.
- Building a path from reaction type to stoichiometry, net ionic equations and limiting reagent calculations.
- Explaining why two equations that look similar may use different chemistry rules.
Common mistakes
- Calling every reaction with oxygen combustion even when CO2 and H2O are not the products.
- Confusing double replacement with single replacement because both may contain ionic compounds.
- Classifying from coefficients instead of the reactant-product pattern.
- Assuming a predicted type proves the reaction occurs under all conditions.
Result checking
After the type appears, compare it with the balanced equation and the number of substances on each side. Then use a second rule set when needed: activity series for metals, solubility rules for precipitates, and acid-base ideas for neutralization.
Related Chemistry Tools
FAQs
What is a combustion reaction in chemistry?
A combustion reaction is one in which a substance reacts rapidly with oxygen to release heat and (usually) light. Complete combustion of a hydrocarbon yields CO2 and H2O — for example, CH4 + 2O2 → CO2 + 2H2O, ΔH ≈ -890 kJ/mol. Incomplete combustion, with limited O2, produces CO and/or soot: 2CH4 + 3O2 → 2CO + 4H2O. Other examples: 2C8H18 + 25O2 → 16CO2 + 18H2O (octane); 2Mg + O2 → 2MgO; C2H5OH + 3O2 → 2CO2 + 3H2O. Combustion reactions are exothermic, fast, and require fuel, oxidizer, and an ignition source.
What is combustion in the carbon cycle?
Combustion is one of the four main processes that move carbon between reservoirs (atmosphere, oceans, biosphere, lithosphere). Burning organic matter or fossil fuels releases stored carbon back to the atmosphere as CO2: C + O2 → CO2; CH4 + 2O2 → CO2 + 2H2O; (C6H10O5)n + n·6O2 → n·6CO2 + n·5H2O. Burning recently grown biomass (wood) returns carbon that was photosynthesized recently. Burning fossil fuels (coal, oil, gas) adds carbon that was sequestered for millions of years, raising atmospheric CO2 from a pre-industrial 280 ppm to over 420 ppm today. Photosynthesis (6CO2 + 6H2O → C6H12O6 + 6O2) is the reverse direction of the same balance.
What are the products of the combustion of a hydrocarbon?
Complete combustion of a hydrocarbon yields CO2 and H2O. General equation: CxHy + (x + y/4) O2 → x CO2 + (y/2) H2O. Examples: CH4 + 2O2 → CO2 + 2H2O; C3H8 + 5O2 → 3CO2 + 4H2O; 2C8H18 + 25O2 → 16CO2 + 18H2O. If oxygen is limited, combustion is incomplete and the products include carbon monoxide (CO) and/or unburned carbon as soot — for example, 2CH4 + 3O2 → 2CO + 4H2O. Visual cue on a gas stove: a steady blue flame indicates complete combustion; a yellow or orange flame typically means incomplete combustion and CO production.
Does combustion release carbon dioxide?
Yes, when the fuel contains carbon. All hydrocarbon and carbohydrate combustion produces CO2 if the reaction goes to completion: CH4 + 2O2 → CO2 + 2H2O; 2C8H18 + 25O2 → 16CO2 + 18H2O; C6H12O6 + 6O2 → 6CO2 + 6H2O. Combustion of hydrogen gas produces only water (2H2 + O2 → 2H2O) — no CO2 — which is why H2 is being explored as a clean energy carrier. Metal combustion (for example, 2Mg + O2 → 2MgO) produces a metal oxide, again with no CO2. Incomplete combustion of carbon-containing fuels produces CO instead of (or alongside) CO2.
What is a synthesis chemical reaction?
A synthesis (combination) reaction has two or more reactants combining into a single product: A + B → AB. Element + element: 2H2 + O2 → 2H2O; N2 + 3H2 → 2NH3 (Haber-Bosch); 2Na + Cl2 → 2NaCl; 2Mg + O2 → 2MgO. Compound + compound: CaO + H2O → Ca(OH)2 (slaking of lime); SO3 + H2O → H2SO4; NH3 + HCl → NH4Cl. Element + compound: 2SO2 + O2 → 2SO3 (Contact process step). Most synthesis reactions are exothermic and form the backbone of industrial chemistry — ammonia, sulfuric acid, and cement production all rely on synthesis steps.
What is decomposition in chemistry?
A decomposition reaction is the reverse of synthesis: a single compound breaks into two or more products. General form: AB → A + B. Three common drivers: (1) Heat (thermal): CaCO3 → CaO + CO2; 2KClO3 → 2KCl + 3O2 (with MnO2 catalyst); 2NaHCO3 → Na2CO3 + H2O + CO2. (2) Electricity (electrolytic): 2H2O → 2H2 + O2; 2NaCl(l) → 2Na + Cl2 (Down's cell); 2Al2O3 → 4Al + 3O2 (Hall-Héroult process). (3) Light (photolytic): 2AgBr → 2Ag + Br2 (silver halide photography); 2H2O2 → 2H2O + O2 in sunlight. Most decomposition reactions are endothermic.
What type of substances are used in decomposition?
Common substances that decompose under heat, light, or electricity include: (1) Carbonates and bicarbonates: CaCO3 → CaO + CO2; 2NaHCO3 → Na2CO3 + H2O + CO2. (2) Metal hydroxides: Cu(OH)2 → CuO + H2O; 2Fe(OH)3 → Fe2O3 + 3H2O. (3) Nitrates: 2KNO3 → 2KNO2 + O2 (alkali metal nitrates); 2Cu(NO3)2 → 2CuO + 4NO2 + O2 (heavy metal nitrates). (4) Oxides of less reactive metals: 2HgO → 2Hg + O2; 2Ag2O → 4Ag + O2. (5) Peroxides: 2H2O2 → 2H2O + O2 (with MnO2 or catalase). (6) Hydrates losing water: CuSO4·5H2O → CuSO4 + 5H2O. (7) Chlorates: 2KClO3 → 2KCl + 3O2. (8) Silver halides under light: 2AgCl → 2Ag + Cl2.
What is the conjugate base?
The conjugate base is what is left after an acid donates a proton (Brønsted-Lowry definition). If HA is the acid, its conjugate base is A-. Examples: HCl → Cl-; CH3COOH → CH3COO-; NH4+ → NH3; H2O → OH-; H2SO4 → HSO4-. Strength is inversely paired: the stronger the acid, the weaker its conjugate base, and vice versa. HCl is strong, so Cl- is a very weak base; acetic acid is weak, so acetate is a moderately strong base. Buffer solutions exploit this pairing — a weak acid with its conjugate base resists pH change. The acetate buffer (CH3COOH / CH3COO-) buffers around pH 4.7; blood is buffered by the H2CO3 / HCO3- pair near pH 7.4.
What is the difference between an acid and a base?
Three theories define acids and bases at increasing levels of generality. Arrhenius: an acid releases H+ in water; a base releases OH-. Brønsted-Lowry: an acid is a proton donor; a base is a proton acceptor. Lewis: an acid accepts an electron pair; a base donates one. Quick contrasts: acids taste sour, turn blue litmus red, react with active metals to release H2, and have pH below 7; bases taste bitter, feel slippery, turn red litmus blue, and have pH above 7. Neutralization brings them together: acid + base → salt + water (e.g., HCl + NaOH → NaCl + H2O).
How to tell if an acid is strong or weak?
Strong acids ionize essentially completely in water; weak acids ionize only partially. There are seven common strong acids worth memorizing: HCl, HBr, HI, HNO3, H2SO4 (first dissociation), HClO3, and HClO4. Anything else in introductory chemistry is weak. Quantitative checks: a large Ka (much greater than 1) or a negative pKa indicates a strong acid; a small Ka (typically 10^-3 to 10^-10) and a positive pKa indicate a weak acid. For 0.1 M solutions, pH is close to 1 for a strong acid (HCl) but around 2.9 for a weak one (acetic acid). Strong acids also conduct electricity better and react more vigorously with active metals.
What happens when you mix an acid and a base?
An acid and a base react in a neutralization: acid + base → salt + water. The net ionic equation when both are strong is H+ + OH- → H2O. Examples: HCl + NaOH → NaCl + H2O (resulting pH = 7); 2HCl + Ca(OH)2 → CaCl2 + 2H2O; H2SO4 + 2NaOH → Na2SO4 + 2H2O. When acid and base are not equally strong, the salt hydrolyzes and shifts the final pH: weak acid + strong base (CH3COOH + NaOH → CH3COONa + H2O) gives a basic solution; strong acid + weak base (HCl + NH3 → NH4Cl) gives an acidic solution. Neutralization is exothermic — about -57 kJ per mole of water formed for strong acid/strong base in dilute solution.
Is NH3 a weak acid?
No — NH3 is a weak base, not a weak acid. In water, ammonia accepts a proton: NH3 + H2O ⇌ NH4+ + OH-, with Kb ≈ 1.8 × 10^-5 at 25 °C. The lone pair on nitrogen accepts H+. NH3 can act as an acid in extreme conditions — for example, the self-ionization of liquid ammonia gives NH4+ and NH2- — but in aqueous chemistry it behaves as a base. The conjugate acid of NH3 is the ammonium ion, NH4+, which is itself a weak acid (Ka ≈ 5.6 × 10^-10). That is why ammonium chloride solutions are slightly acidic.
Are acids slippery?
No — slipperiness is a hallmark of strong bases, not acids. Acids have a sour taste, sting on contact, turn blue litmus red, react with active metals to release H2, and react with carbonates to release CO2. Bases like NaOH and KOH feel slippery because they saponify oils and fats on the skin — converting them into soap. That slippery feeling is actually skin damage in progress, so concentrated bases need careful handling. Other slippery bases include ammonia solution and lime water (Ca(OH)2). If you contact a strong base, rinse with plenty of water immediately.
Are all acids aqueous?
No — acids exist in every physical state. Gaseous: HCl gas, H2S, and acidic oxides such as SO2, NO2, and CO2. Liquid: pure H2SO4 (a viscous liquid above 10 °C), HNO3, and HF. Solid: citric acid, tartaric acid, oxalic acid, and boric acid. Acids only need water to ionize as Brønsted acids; Lewis acids such as BF3, AlCl3, and FeCl3 act as electron-pair acceptors without any solvent or proton. Aqueous solutions are simply the most common laboratory form because ionization in water is convenient and well-characterized.
Do all acids have hydrogen?
Most acids contain hydrogen, but not all. Arrhenius and Brønsted-Lowry acids must contain ionizable hydrogen (HCl, H2SO4, CH3COOH, etc.), since both definitions involve donating H+. Lewis acids only need to accept an electron pair, so many contain no hydrogen at all: BF3, AlCl3, FeCl3, SO3, and metal cations such as Al3+ and Cu2+ are all Lewis acids. BF3 + NH3 → F3B-NH3 illustrates a Lewis acid-base reaction with no proton transfer. Acidic oxides (SO3, CO2, P2O5) become Brønsted acids only after reacting with water — for example, SO3 + H2O → H2SO4.
Is H2O a base?
Water is amphoteric — it can act as either an acid or a base depending on its partner. As a base: HCl + H2O → H3O+ + Cl- (water accepts a proton). As an acid: NH3 + H2O ⇌ NH4+ + OH- (water donates a proton). Pure water also self-ionizes: 2H2O ⇌ H3O+ + OH-, with Kw = 1.0 × 10^-14 at 25 °C. The two lone pairs on oxygen let it accept a proton; the two O-H bonds let it donate one. This dual behavior is why water is the universal reaction medium for acid-base chemistry.