Difficulty: Introductory | Prerequisites: familiarity with chemical formulas, states of matter notation (s, l, g, aq).
This lab covers the five major categories of chemical reactions and how to recognise that a reaction has occurred. It sits early in a general chemistry sequence, right after learning to write and balance equations. If you can name elements and write simple formulas, you are ready. Everything here feeds directly into stoichiometry, redox chemistry, and solution chemistry later in the course.
Chemical reactions fall into five main types: combination, decomposition, single replacement, double replacement, and combustion. You can tell a reaction has occurred by watching for colour changes, gas bubbles, precipitate formation, or temperature changes. Learning to classify reactions by their pattern lets you predict products and balance equations, which is most of what this part of the course tests.
Combination reaction (synthesis)
A reaction in which two or more substances combine to form a single product. Think of it as: two things joining to make one new thing. A + B → AB.
Decomposition reaction
A reaction in which a single compound breaks down into two or more simpler substances. In simple terms, this means one thing falls apart into smaller pieces. AB → A + B.
Single replacement reaction (single displacement)
A reaction in which one element replaces another element in a compound. Think of it as: one element cutting in line and kicking the other one out. A + BC → AC + B.
Double replacement reaction (double displacement, metathesis)
A reaction in which the positive ions of two ionic compounds swap partners. In simple terms, two compounds trade parts with each other. AB + CD → AD + CB.
Combustion reaction
A reaction in which a hydrocarbon (or other fuel) reacts with oxygen to produce carbon dioxide and water, releasing energy. Think of it as: burning. Fuel + O₂ → CO₂ + H₂O.
Activity series
A ranked list of metals ordered by their tendency to lose electrons and react. A metal higher on the list can displace a metal lower on the list from a compound. In simple terms, it tells you which metals are "stronger" and can push "weaker" metals out of solution. That is why zinc displaces copper from CuSO₄, but copper cannot displace zinc.
Precipitate
An insoluble solid that forms when two aqueous solutions are mixed. Think of it as: the solid that crashes out of a clear liquid when you mix two solutions together, like the white Ca₃(PO₄)₂ that appeared when CaCl₂ and Na₃PO₄ were combined.
Balanced equation
A chemical equation in which the number of atoms of each element is the same on both sides of the arrow. In simple terms, what goes in must come out. You adjust coefficients (the big numbers in front of formulas) until every atom balances.
Evidence of a chemical reaction
Observable signs that a chemical change has taken place, including colour change, gas production (bubbles), precipitate formation, temperature change, or emission of light. Think of it as: the clues that tell you something new has been made, rather than just mixed.
Two or more reactants combine into one product.
Lab example: burning magnesium ribbon in air.
Initial appearance: shiny, light grey ribbon of metal.
Evidence: intense heat and white light produced; white ash (MgO) left behind.
Balanced equation: 2 Mg(s) + O₂(g) → 2 MgO(s)
Other examples from the problem set:
Ni(s) + F₂(g) → NiF₂(s)
H₂(g) + Cl₂(g) → 2 HCl(aq)
4 Fe(s) + 3 O₂(g) → 2 Fe₂O₃(s)
A single compound breaks into two or more simpler substances.
Lab example: hydrogen peroxide with potassium iodide (KI) catalyst.
Appearance: clear, colourless liquid.
Evidence: formation of bubbles (oxygen gas released).
Balanced equation: 2 H₂O₂(l) → 2 H₂O(l) + O₂(g)
Other examples:
CaCO₃(s) → CaO(s) + CO₂(g)
MgCO₃(s) → MgO(s) + CO₂(g)
Sodium carbonate + HCl is classified in this lab as a double replacement followed by decomposition, because the initial product (H₂CO₃) immediately decomposes into CO₂ and H₂O.
One element displaces another from a compound. Requires the free element to be higher in the activity series than the element it replaces.
Lab examples:
Zinc in copper(II) sulfate solution:
CuSO₄(aq) started blue/green; Zn(s) was dull grey.
After 30 min: solution turned light blue; zinc turned black (copper deposited on it).
Balanced equation: Zn(s) + CuSO₄(aq) → Cu(s) + ZnSO₄(aq)
Metals in hydrochloric acid:
Zn(s) + 2 HCl(aq) → ZnCl₂(aq) + H₂(g) — bubbling observed.
Cu(s) + HCl(aq) → NR (no reaction). Copper is below hydrogen in the activity series.
Mg(s) + 2 HCl(aq) → MgCl₂(aq) + H₂(g) — vigorous bubbling and heat.
Other examples:
2 Zn(s) + CuBr₂(aq) → 2 ZnBr(aq) + Cu(s)
Fe₂O₃(s) + 3 C(s) → 2 Fe(s) + 3 CO₂(g) (reduction of iron ore)
The cations of two ionic compounds swap anions. Typically occurs in aqueous solution and is driven by formation of a precipitate, water, or a gas.
Lab examples:
CaCl₂(aq) + Na₃PO₄(aq):
Both solutions clear and colourless.
Evidence: white solid (precipitate) formed.
Balanced equation: 3 CaCl₂(aq) + 2 Na₃PO₄(aq) → Ca₃(PO₄)₂(s) + 6 NaCl(aq)
FeCl₃(aq) + KSCN(aq):
FeCl₃ was orange; KSCN was clear and colourless.
Evidence: colour change to blood red (formation of Fe(SCN)₃).
Balanced equation: FeCl₃(aq) + 3 KSCN(aq) → Fe(SCN)₃(aq) + 3 KCl(aq)
Na₂CO₃(s) + HCl(aq):
Evidence: vigorous bubbling (CO₂ gas). A burning splint went out due to CO₂ displacing oxygen.
Balanced equation: Na₂CO₃(s) + 2 HCl(aq) → CO₂(g) + H₂O(l) + 2 NaCl(aq)
Other examples:
Ca(OH)₂(aq) + 2 HNO₃(aq) → Ca(NO₃)₂(aq) + 2 H₂O(l) (acid-base neutralisation)
KCl(aq) + AgNO₃(aq) → KNO₃(aq) + AgCl(s)
A substance (usually a hydrocarbon) reacts with O₂ to produce CO₂ and H₂O, releasing heat and light.
From the problem set:
C₅H₁₂(l) + 8 O₂(g) → 5 CO₂(g) + 6 H₂O(g) (pentane burning)
C₂H₄(g) + 3 O₂(g) → 2 CO₂(g) + 2 H₂O(g)
Always look for a hydrocarbon + O₂ on the reactant side. Products are always CO₂ + H₂O for complete combustion.
Reaction Type | General Form | How to Spot It |
|---|---|---|
Combination | A + B → AB | Two or more reactants, one product |
Decomposition | AB → A + B | One reactant, two or more products |
Single replacement | A + BC → AC + B | An element and a compound on the left; a different element freed on the right |
Double replacement | AB + CD → AD + CB | Two compounds swap ions; look for a precipitate, water, or gas as a driving force |
Combustion | CₓHₙ + O₂ → CO₂ + H₂O | A hydrocarbon reacts with oxygen; heat and light released |
Combination: rusting of iron (4 Fe + 3 O₂ → 2 Fe₂O₃) is why bridges and cars corrode over time. Protective coatings exist to block this reaction.
Decomposition: hydrogen peroxide decomposing into water and oxygen is why your bottle of H₂O₂ slowly loses potency on the shelf, and why it fizzes when applied to a wound (catalase in your blood acts like the KI catalyst).
Single replacement: galvanising steel with zinc works because zinc is more reactive than iron, so it corrodes preferentially and protects the steel underneath.
Double replacement: water treatment plants use double replacement reactions to remove dissolved metals as insoluble precipitates.
Combustion: every internal combustion engine burns hydrocarbons (petrol, diesel) with oxygen. The balanced equation predicts how much CO₂ a car produces per litre of fuel.
Students often think that bubbling always means a reaction occurred. It does not. Boiling water bubbles too. The bubbles must be gas produced by the reaction itself (e.g. H₂ or CO₂), not dissolved air escaping.
Students frequently confuse single and double replacement. Count the reactants: if one is a free element, it is single replacement. If both reactants are compounds (ionic or acid-base), it is double replacement.
A common error is assuming copper reacts with HCl. Copper sits below hydrogen in the activity series, so Cu + HCl → NR. The activity series is not optional knowledge here.
Students sometimes write combustion products as CO instead of CO₂. Unless the question specifies incomplete combustion, assume complete combustion and write CO₂ + H₂O.
⚠️ Classifying a reaction by type is one of the most common exam questions. Learn the five general forms cold.
⚠️ Expect questions that give you reactants and ask you to predict products and balance. The type tells you the product pattern.
⚠️ The activity series determines whether a single replacement reaction proceeds or yields NR (no reaction). Know at minimum that Mg > Zn > Fe > H > Cu.
⚠️ Na₂CO₃ + HCl is a favourite test question because it is a double replacement that immediately decomposes, so students must identify both steps.
⚠️ Evidence of a chemical reaction is frequently tested as a multiple-choice question: colour change, gas evolution, precipitate formation, temperature change, light emission.
True or False: A single replacement reaction has two compounds as reactants. (False. One reactant is a free element.)
Fill in the blank: In a combustion reaction, a hydrocarbon reacts with ______ to produce CO₂ and H₂O. (oxygen / O₂)
True or False: Copper reacts with hydrochloric acid to produce hydrogen gas. (False. Cu is below H in the activity series, so no reaction occurs.)
Fill in the blank: When two clear aqueous solutions are mixed and a solid appears, that solid is called a ______. (precipitate)
True or False: Decomposition reactions have two or more reactants. (False. They have one reactant that breaks into two or more products.)
Q: Classify the following reaction: 2 Na(s) + 2 H₂O(l) → 2 NaOH(aq) + H₂(g)
A: Single replacement. Sodium (a free element) displaces hydrogen from water.
Q: What evidence of a chemical reaction would you expect when burning a match?
A: Heat and light (flame), along with a colour change as the match head transforms into ash.
Q: Write and balance the equation for the combustion of pentane (C₅H₁₂).
A: C₅H₁₂(l) + 8 O₂(g) → 5 CO₂(g) + 6 H₂O(g)
Q: Why does copper not react with hydrochloric acid, while zinc does?
A: Copper is below hydrogen in the activity series, so it cannot displace hydrogen from HCl. Zinc is above hydrogen, so it displaces H₂ gas.
Q: When CaCl₂(aq) is mixed with Na₃PO₄(aq), a white solid forms. Write the balanced equation, name the type of reaction, and identify the precipitate.
A: 3 CaCl₂(aq) + 2 Na₃PO₄(aq) → Ca₃(PO₄)₂(s) + 6 NaCl(aq). Double replacement. The precipitate is calcium phosphate, Ca₃(PO₄)₂.
Q: A burning splint is held over a test tube where Na₂CO₃ is reacting with HCl. The flame goes out. Explain why.
A: The reaction produces CO₂ gas, which is denser than air and not flammable. It displaces the oxygen around the splint, extinguishing the flame.
Q: Classify CaCO₃(s) → CaO(s) + CO₂(g).
A: Decomposition. One compound breaks into two simpler substances.
This material connects directly to stoichiometry, because once you can predict products and balance equations you can calculate masses and moles of reactants and products. It also feeds into redox chemistry: single replacement reactions are a subset of redox, and the activity series is essentially a ranking of reduction potentials. Double replacement reactions reappear in the solubility rules and net ionic equations unit, where you learn to identify spectator ions and write net ionic equations for precipitate-forming reactions.
chemical reactions, types of reactions, combination reaction, synthesis reaction, decomposition reaction, single replacement, single displacement, double replacement, double displacement, metathesis, combustion, activity series, reactivity series, precipitate, balancing equations, evidence of chemical reaction, Lab 10, general chemistry, CHM, Michigan State, stoichiometry, redox, oxidation-reduction, net ionic equation, hydrocarbon combustion, NR no reaction