Lesson 5

Predicting and Balancing Chemical Reactions

Balance chemical equations and predict products across the five major reaction families: synthesis, decomposition, single replacement, double replacement, and combustion.

6 learning objectivesreactions
Chemistry reference tables

Why Equations Must Be Balanced

A chemical equation is a quantitative statement: it says exactly how many atoms of each element take part in the reaction. The driving idea is the law of conservation of mass — in any chemical change, atoms are neither created nor destroyed. Every atom on the reactant side has to appear on the product side. The same is true for charge in ionic equations.

Why this matters in practice:

  • Stoichiometry depends on it. The coefficients in a balanced equation are mole ratios. An unbalanced equation gives the wrong moles, the wrong masses, the wrong yield — everything downstream breaks.
  • It encodes the actual chemistry. 2 H2 + O2 → 2 H2O is a real claim about how hydrogen and oxygen combine. Writing H2 + O2 → H2O implicitly destroys one oxygen atom and tells a chemical lie.
  • It exposes errors in the predicted products. If a proposed product set cannot be balanced (for instance, you wrote the wrong charge on an ion and ended up with an impossible atom count), that is a tell that the products themselves are wrong. Balancing is a sanity check on prediction.

Two rules that follow from atom conservation: subscripts in formulas are fixed (the formula identifies the substance — you cannot change H2O to H2O2 to balance an equation), but coefficients in front of formulas are adjustable. Balancing is exactly the process of choosing those coefficients so atoms (and charge) balance on both sides. By convention, the smallest whole-number set of coefficients is preferred.

Go deeperWhy this matters: where fat actually goes

When someone loses 10 kg of body fat, where does it go? Surveys have found that most people, including many doctors, guess “energy,” “heat,” or “sweat.” Conservation of mass says otherwise: the carbon and hydrogen atoms in fat must leave as atoms. The arithmetic works out to roughly 8.4 kg exhaled as CO⁠2 and 1.6 kg leaving as water. You breathe out most of your lost weight. Every balanced equation you write is this same bookkeeping: atoms in, atoms out, nothing vanishes.

Go deeperThe sealed flasks that proved the law

Burning wood seems to destroy mass: the ash weighs far less than the log. Antoine Lavoisier settled the question in the 1770s by running reactions in sealed vessels and weighing everything, gases included, before and after. The totals always matched. Mass only appears to vanish when a gaseous product escapes uncounted, and it only appears to grow (as when iron rusts and gains weight) when a gas is quietly consumed from the air.

That is why phase labels like (g) deserve respect in your equations: the gases are where the “missing” mass hides.

Balancing Simple Molecular Equations

Balancing a molecular equation is a systematic counting exercise. A reliable workflow:

  1. Write the unbalanced ("skeleton") equation with correct formulas for every reactant and product, including phase labels (s), (l), (g), (aq) when given.
  2. Count atoms of each element on both sides.
  3. Pick the most complex molecule first (often the one with the most distinct elements) and place a coefficient that balances one of its elements. Save free elements (H2, O2, metals) for last.
  4. Cycle through elements: balance metals, then non-metals other than H and O, then H, then O. Update counts after every change.
  5. Reduce coefficients to their lowest whole-number ratio at the end.
  6. Verify: every element has the same total atom count on both sides.

Worked example. Balance C3H8 + O2 → CO2 + H2O.

  1. Carbon: 3 on the left, 1 on the right. Place a 3 on CO2: C3H8 + O2 → 3 CO2 + H2O.
  2. Hydrogen: 8 on the left, 2 on the right. Place a 4 on H2O: C3H8 + O2 → 3 CO2 + 4 H2O.
  3. Oxygen: 2 on the left, (3 × 2) + (4 × 1) = 10 on the right. Place a 5 on O2: C3H8 + 5 O2 → 3 CO2 + 4 H2O.
  4. Check: C 3=3, H 8=8, O 10=10. Balanced.

Two common pitfalls. First, do not change subscripts — that changes the substance. Second, when an element appears in multiple compounds on the same side, balance it last (this is why H and O usually go after the metals/nonmetals). For odd–even mismatches involving free O2, you can use a fractional coefficient ( + ½ O2) and then double everything at the end to clear the fraction.

Go deeperCommon mistake: fixing the count by changing the formula

Faced with H⁠2 + O⁠2 → H⁠2O, it is tempting to “balance” the oxygens by writing H⁠2O⁠2 on the right. The atom counts now match, but you have changed the product from water to hydrogen peroxide, a different substance entirely. The reaction no longer describes what actually happens.

The rule is absolute: subscripts are fixed by the identity of each substance; only coefficients may change. Coefficients say how much; subscripts say what. The correct balance is 2 H⁠2 + O⁠2 → 2 H⁠2O, achieved by adjusting amounts, never identities.

Go deeperThe fraction trick for stubborn oxygens

Combustion equations often leave you needing an odd number of oxygen atoms from even-numbered O⁠2 molecules. Allow yourself a temporary fraction: C⁠2H⁠6 + 7/2 O⁠2 → 2 CO⁠2 + 3 H⁠2O balances perfectly (seven O atoms needed on the left, and 7/2 molecules of O⁠2 supply them).

Then clear the fraction by doubling every coefficient: 2 C⁠2H⁠6 + 7 O⁠2 → 4 CO⁠2 + 6 H⁠2O. The fraction was scaffolding, not a final answer: an equation read at the molecular level cannot have half a molecule, but as an intermediate step the fraction turns a frustrating trial-and-error hunt into two quick moves.

Classifying Chemical Reactions by Type

Chemical reactions can be organized into five major categories. Recognizing the type helps you predict the products:

  1. Synthesis (combination): Two or more substances combine to form one product. General form: A + B → AB.
  2. Decomposition: One compound breaks down into two or more simpler substances. General form: AB → A + B.
  3. Single replacement (displacement): An element replaces another element in a compound. General form: A + BC → AC + B.
  4. Double replacement (metathesis): Two compounds exchange ions to form two new compounds. General form: AB + CD → AD + CB.
  5. Combustion: A substance reacts with oxygen, producing heat and light. Hydrocarbons burn to produce CO2 and H2O.

To classify a reaction: count the number of reactants and products, and check whether elements or compounds are exchanging partners. When a reaction occurs in aqueous solution and the driving force is precipitation, acid-base neutralization, or gas formation, the aqueous-specific patterns covered in Topic 6 take precedence over the general categories here.

Go deeperTest yourself: classify three reactions

Name the reaction type for each, then check below.

  • 2 KClO⁠3 → 2 KCl + 3 O⁠2
  • Zn + CuSO⁠4 → ZnSO⁠4 + Cu
  • C⁠3H⁠8 + 5 O⁠2 → 3 CO⁠2 + 4 H⁠2O

Answers: The first is decomposition (one compound breaking into simpler substances). The second is single replacement (free zinc displaces copper from its compound). The third is combustion (a fuel plus O⁠2 giving CO⁠2 and H⁠2O). Notice how the shape of the equation gives the type away before you check a single atom: one reactant, element-plus-compound, or fuel-plus-oxygen.

Synthesis Reactions

In a synthesis reaction (also called a combination reaction), two or more reactants combine to form a single, more complex product. The general form is A + B → AB. Common synthesis patterns:

  • Metal + nonmetal → expected binary product formula for the introductory pairs used here: use the common ion charges to write the product formula, as in 2 Na(s) + Cl⁠2(g) → 2 NaCl(s) and 2 Mg(s) + O⁠2(g) → 2 MgO(s). This is a product-formula heuristic, not a universal bonding classification; anhydrous AlBr⁠3, for example, is covalent.
  • Nonmetal + nonmetal → covalent or hybrid compound: H⁠2(g) + Cl⁠2(g) → 2 HCl(g); N⁠2(g) + 3 H⁠2(g) → 2 NH⁠3(g) (Haber process); C(s) + O⁠2(g) → CO⁠2(g). When the two reactants are a metal and H⁠2, the resulting hydride is ionic: 2 Na(s) + H⁠2(g) → 2 NaH(s).
  • Metal oxide + water → metal hydroxide: Na⁠2O(s) + H⁠2O(l) → 2 NaOH(aq); CaO(s) + H⁠2O(l) → Ca(OH)⁠2(aq). Only soluble metal oxides (mostly Group 1 and the heavier Group 2 oxides) do this readily.
  • Nonmetal oxide + water → oxyacid: SO⁠3(g) + H⁠2O(l) → H⁠2SO⁠4(aq); Cl⁠2O⁠7(g) + H⁠2O(l) → 2 HClO⁠4(aq); CO⁠2(g) + H⁠2O(l) → H⁠2CO⁠3(aq). Nonmetal oxides act as “acid anhydrides.”
  • Metal oxide + nonmetal oxide → salt: CaO(s) + CO⁠2(g) → CaCO⁠3(s); CaO(s) + SO⁠3(g) → CaSO⁠4(s). The metal oxide donates O to form the polyatomic anion paired with the metal cation.
  • Lower oxide + O⁠2 → higher oxide (further oxidation): P⁠4O⁠6(s) + 2 O⁠2(g) → P⁠4O⁠10(s); 2 SO⁠2(g) + O⁠2(g) → 2 SO⁠3(g) (industrially important).

To predict a synthesis product, recognize the reactant categories, match to one of the patterns above, and then write the correct product formula using ionic charges or known molecular formulas.

Go deeperWhy this matters: the synthesis reaction that feeds the world

The unassuming equation N⁠2 + 3 H⁠2 → 2 NH⁠3 is arguably the most consequential chemical reaction of the last century. The Haber process fixes inert atmospheric nitrogen into ammonia for fertilizer, and by widely cited estimates, roughly 40 to 50% of the nitrogen atoms in your body’s proteins have passed through an industrial Haber reactor. Running it consumes on the order of 1 to 2% of the world’s energy supply.

Without this one synthesis reaction, Earth’s farmland could feed only about half its current population. Product prediction is not just exam technique; sometimes the predicted product changes history.

Common Product-Prediction Patterns (LO 5.4)

LO 5.4 asks you to predict products of common reactions across categories that go beyond just synthesis. The patterns below cross several reaction types — some are redox single-replacement, some are acid-base, some are double-replacement — but they share a common goal: given a metal/nonmetal/oxide/hydroxide/acid/salt reactant pair, predict the products.

  • Active metal + water → metal hydroxide + H⁠2(g) (single-replacement / redox, not synthesis): Group 1 metals (Li, Na, K, Rb, Cs) react vigorously with water; reactivity increases down the group: 2 Cs(s) + 2 H⁠2O(l) → 2 CsOH(aq) + H⁠2(g). Heavier Group 2 metals (Ca, Sr, Ba) react more slowly: Ca(s) + 2 H⁠2O(l) → Ca(OH)⁠2(aq) + H⁠2(g). Group 2 is not uniform: beryllium does not react with water at all, and magnesium reacts only very slowly with cold water (readily with steam, giving MgO + H⁠2). Aluminium is protected by a tight oxide layer (passivation) and does not react with water under ordinary conditions, even though it sits above hydrogen in the activity series.
  • Metal oxide + acid → salt + water (acid-base, not synthesis): Fe⁠2O⁠3(s) + 6 HCl(aq) → 2 FeCl⁠3(aq) + 3 H⁠2O(l); CuO(s) + H⁠2SO⁠4(aq) → CuSO⁠4(aq) + H⁠2O(l). Metal oxides are “basic” toward acids.
  • Nonmetal oxide + base → salt + water (acid-base, not synthesis): CO⁠2(g) + 2 NaOH(aq) → Na⁠2CO⁠3(aq) + H⁠2O(l); SO⁠2(g) + 2 KOH(aq) → K⁠2SO⁠3(aq) + H⁠2O(l). Nonmetal oxides are “acidic” toward bases.
  • Acid + salt → new acid + new salt (double-replacement, volatility-driven, not synthesis): a non-volatile concentrated acid can displace a volatile acid from its salt. H⁠2SO⁠4(l) + NaCl(s) → NaHSO⁠4(s) + HCl(g) is the classic preparation of HCl gas; the volatile HCl escapes, driving the reaction forward.
  • Halide + halogen → higher halide (single-replacement / redox, not synthesis): a more electronegative halogen oxidizes the metal in a halide salt to a higher oxidation state. 2 FeCl⁠2(s) + Cl⁠2(g) → 2 FeCl⁠3(s).

Treat each pattern as: identify the reactant categories, look up which pattern matches, write the products with correct formulas. The reaction-type classification (synthesis vs single-replacement vs double-replacement vs acid-base) for each pattern is noted explicitly so you do not blur LO 5.3 (classification) with LO 5.4 (product prediction).

Decomposition Reactions

In a decomposition reaction, a single compound breaks apart into two or more simpler substances, often when heated. Common decomposition patterns:

  • Metal carbonates decompose into metal oxide + CO⁠2: CaCO⁠3(s) → CaO(s) + CO⁠2(g)
  • Metal hydroxides decompose into metal oxide + H⁠2O: Ca(OH)⁠2(s) → CaO(s) + H⁠2O(g)
  • Metal chlorates decompose into metal chloride + O⁠2: 2 KClO⁠3(s) → 2 KCl(s) + 3 O⁠2(g)
  • Binary compounds decompose into elements: 2 H⁠2O(l) → 2 H⁠2(g) + O⁠2(g)

Decomposition reactions are often the reverse of synthesis reactions. Heat, electricity, or light typically provide the energy needed to break chemical bonds.

Go deeperWhy this matters: the decomposition that saves your life in 30 milliseconds

A classic airbag inflator is a decomposition reaction on a hair trigger: a spark ignites sodium azide, 2 NaN⁠3 → 2 Na + 3 N⁠2, flooding the bag with nitrogen gas faster than your head can reach the steering wheel. The same logic, gentler and slower, raises your muffins: baking soda decomposes when heated, releasing the CO⁠2 that makes batter rise.

Both applications exploit the defining feature of this reaction class: one stable solid, triggered on demand, releasing a large volume of gas.

Go deeperTry it: predict the products of heating magnesium carbonate

Using the carbonate pattern, predict what MgCO⁠3(s) becomes when strongly heated, then check below.

Answer: MgCO⁠3(s) → MgO(s) + CO⁠2(g), following the pattern metal carbonate → metal oxide + carbon dioxide. The equation is already balanced as written (one Mg, one C, three O on each side). This same pattern, run on limestone (CaCO⁠3) in kilns at industrial scale, produces the lime used in cement, and it is why cement manufacturing releases CO⁠2 even before any fuel is burned.

Single Replacement Reactions and the Activity Series

In a single replacement reaction, a free element displaces another element from a compound. Whether the reaction occurs depends on the activity series — a ranking of elements by their tendency to be oxidized.

The rule: A more active element can displace a less active element from a compound, but not vice versa.

For metals, a partial activity series (most active first): Li > K > Ca > Na > Mg > Al > Mn > Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au.

  • Zn + CuSO⁠4 → ZnSO⁠4 + Cu (Zn is above Cu, so reaction occurs)
  • Cu + ZnSO⁠4 → no reaction (Cu is below Zn)
  • Metals above H in the series can displace hydrogen from acids: Zn + 2 HCl → ZnCl⁠2 + H⁠2

For halogens, activity decreases down the group: F⁠2 > Cl⁠2 > Br⁠2 > I⁠2. A more active halogen can displace a less active one from a solution of its salt.

Go deeperWhy this matters: welding rails with a bucket of powder

Railroad crews join track sections with thermite: 2 Al + Fe⁠2O⁠3 → Al⁠2O⁠3 + 2 Fe. Aluminum sits above iron in the activity series, so it rips the oxygen away from iron oxide, and the reaction runs hot enough to deliver molten iron directly into the joint, no power supply required.

The activity series also explains sacrificial protection: zinc blocks bolted to ship hulls and buried alongside steel pipelines corrode first, sparing the iron, because zinc is the more active metal. Engineers read the activity series the way you are learning to: as a table of who displaces whom.

Go deeperTry it: which beaker reacts?

Beaker A holds blue copper(II) sulfate solution with a strip of iron. Beaker B holds pale-green iron(II) sulfate solution with a strip of copper. Which one reacts? Check below.

Answer: Only beaker A: iron is above copper in the activity series, so Fe + CuSO⁠4 → FeSO⁠4 + Cu proceeds. You would see the blue color fade as Cu⁠2+ ions leave solution and reddish copper metal plate onto the iron strip. In beaker B, copper sits below iron and cannot displace it: no reaction, and writing “NR” is the correct answer. The series makes both predictions in seconds. (This is also the everyday reason an iron tool left in copper-salt solution comes out copper-colored.)

Double Replacement Reactions (General Pattern)

In a double replacement reaction, two ionic compounds exchange cation-anion partners. The general form: AB + CD → AD + CB. To predict the products, swap the partners (cation A pairs with anion D; cation C pairs with anion B), then write each new compound using charge balance and the rules from nomenclature.

For a double-replacement reaction to actually proceed, a driving force has to remove one of the products from the reaction mixture — an insoluble precipitate, water from acid-base combination, or a gas that escapes. The aqueous-solution chemistry of those three driving forces (precipitation, acid-base neutralization, gas-forming reactions) is the subject of Topic 6, where solubility rules, ionic equations, and the “no reaction” outcome are covered in depth.

At the level of this topic, focus on (1) recognizing the double-replacement pattern from the reactant types, and (2) writing the swap-partner products with correct formulas. The aqueous-specific yes/no decision and the writing of net ionic equations belong to Topic 6.

Go deeperWhy this matters: kidney stones are a double replacement

A kidney stone is precipitation chemistry in the wrong place: dissolved Ca⁠2+ meets dissolved oxalate ions, the partners swap, and insoluble calcium oxalate crystallizes. The scale coating a kettle or boiler is the same story with carbonate: dissolved Ca⁠2+ and CO⁠3⁠2- combining into solid CaCO⁠3. Water softeners exist precisely to intercept this swap, exchanging Ca⁠2+ for Na⁠+ before it can find a partner to precipitate with. When you predict whether a double replacement “goes,” you are doing the same analysis a nephrologist or a boiler engineer does.

Go deeperCommon mistake: predicting a reaction when nothing happens

Swap the partners of NaCl and KNO⁠3 and you get NaNO⁠3 and KCl: all four compounds are soluble, so mixing the two solutions produces no precipitate, no gas, and no water. Nothing is removed from solution; the four ions simply continue to swim independently. The honest answer is “no reaction.”

Writing the swapped products anyway, as if paper chemistry were real chemistry, is the most common error in this reaction class. Always finish with the driving-force check: a double replacement proceeds only if a precipitate, a gas, or water actually forms.

Combustion Reactions

A combustion reaction occurs when a substance (the fuel) reacts rapidly with oxygen gas (O⁠2), producing heat and usually light. The products are predictable based on the fuel’s composition:

  • Hydrocarbons (CxHy) + O⁠2 → CO⁠2 + H⁠2O. Example: CH⁠4 + 2 O⁠2 → CO⁠2 + 2 H⁠2O.
  • Compounds containing C, H, and O + O⁠2 → CO⁠2 + H⁠2O. Example: C⁠2H⁠5OH + 3 O⁠2 → 2 CO⁠2 + 3 H⁠2O.
  • Metals + O⁠2 → metal oxides. Example: 4 Fe + 3 O⁠2 → 2 Fe⁠2O⁠3 (rusting is slow combustion).

Balancing tip for hydrocarbon combustion: balance C first, then H, then balance O last (since O⁠2 is the only source of oxygen on the reactant side, you can adjust its coefficient to match the total O atoms needed).

Go deeperCommon mistake: assuming combustion is always complete

The tidy pattern fuel + O⁠2 → CO⁠2 + H⁠2O holds only when oxygen is plentiful. Starve the flame of O⁠2 and carbon leaves half-finished: as carbon monoxide, or as unburned soot. That is why a healthy gas-stove flame burns blue (complete combustion) while a yellow, flickering flame is glowing soot announcing incomplete combustion, and why furnaces and water heaters need ventilation: a blocked flue turns a CO⁠2 producer into a CO producer.

On an exam, assume complete combustion unless told otherwise, but know that the assumption is a statement about oxygen supply, not a law.

A Decision Strategy for Predicting Products

When faced with an unfamiliar equation, a systematic approach prevents guesswork:

  1. Count the reactants. One reactant suggests decomposition. Two reactants could be synthesis, single replacement, double replacement, or combustion.
  2. Is O2 a reactant and is the other substance a fuel (contains C and/or H)? If yes, it is a combustion reaction — products are CO2 and H2O.
  3. Are both reactants compounds? If yes, this is the double-replacement pattern: swap the cation/anion partners and write the formulas of the two new compounds. At T5 the goal is recognizing the pattern and writing the products. When the reaction occurs in aqueous solution, the additional question of whether a driving force (precipitate, water, gas, or no reaction) actually makes the reaction proceed is part of Topic 6 (LOs 6.5–6.7) — defer that assessment until then.
  4. Is one reactant a free element and the other a compound? This is a single replacement. Use the activity series to decide whether the reaction proceeds.
  5. Are the reactants two elements or simple substances combining? This is a synthesis. Determine the product’s formula using ionic charges or known molecular formulas.

After predicting products, always balance the equation and verify atom counts on both sides. With practice, this decision tree becomes second nature.

Learning Objectives

After studying this topic, you should be able to:

  1. Explain why a chemical equation must be balanced in terms of conservation of atoms
  2. Balance simple molecular equations
  3. Classify reactions as synthesis, decomposition, single replacement, double replacement, or combustion
  4. Predict products of common reactions of metals, nonmetals, oxides, hydroxides, acids, and salts
  5. Predict products of and write balanced equations for complete combustion of hydrocarbons and other organic compounds
  6. Use the activity series to determine whether a single-replacement reaction will occur and predict its products

How-To Procedure

How to Predict the Products of a Chemical Reaction

  1. Examine the reactants: count how many there are and identify whether they are elements, compounds, or both.
  2. Check for combustion first: if O2 is a reactant and the other substance contains C and/or H, the products are CO2 and H2O.
  3. If one reactant is a free element and the other is a compound, it is single replacement. Consult the activity series to confirm the reaction occurs, then swap the free element with the matching element in the compound.
  4. If both reactants are compounds, it is double replacement. Swap the cation-anion partners and write the formulas of the two new compounds. Stop there at T5 — whether an aqueous double-replacement reaction actually proceeds (driving forces: precipitate, water, gas, or no reaction) is assessed in Topic 6 LOs 6.5–6.7, not here.
  5. If two or more simple substances combine into one product, it is synthesis. Use ionic charges or known molecular formulas to write the product.
  6. If a single compound breaks into simpler substances, it is decomposition. Use common decomposition patterns (carbonates, hydroxides, chlorates) to predict the products.
  7. Balance the equation and verify atom counts on both sides.

Worked Example

Predicting Products and Classifying a Reaction

Problem

When a piece of zinc metal is placed in a solution of copper(II) sulfate (CuSO⁠4), a reaction occurs. Predict the products, classify the reaction type, and write the balanced equation.

Solution
  1. Identify the reactants: Zn(s) is a free metal element; CuSO⁠4(aq) is an ionic compound containing Cu⁠2+ and SO⁠4⁠2- ions.
  2. Recognize the pattern: A free element + a compound → single replacement reaction.
  3. Check the activity series: Zn is above Cu, so Zn can displace Cu⁠2+ from solution. The reaction will occur.
  4. Predict products: Zn replaces Cu⁠2+, forming ZnSO⁠4(aq), and Cu metal is deposited: Zn(s) + CuSO⁠4(aq) → ZnSO⁠4(aq) + Cu(s).
  5. Verify balance: Zn 1=1, Cu 1=1, S 1=1, O 4=4. Balanced as written.
Answer

This is a single replacement reaction. Zn(s) + CuSO⁠4(aq) → ZnSO⁠4(aq) + Cu(s). Zinc is more active than copper, so it displaces copper from solution.

Test Your Understanding

A student places a strip of copper metal into a solution of zinc sulfate (ZnSO4). Will a reaction occur? Explain using the activity series.

Practice Problems

conceptual

Classify each reaction type and predict the products: (a) magnesium + oxygen gas, (b) aluminum + copper(II) chloride solution, (c) propane (C3H8) burning in air.

calculation

Calcium carbonate (CaCO3) decomposes when strongly heated, releasing carbon dioxide gas. Classify the reaction type, predict the second product, and write the balanced equation.

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Self-Study Questions

Why must a chemical equation be balanced?

Hint: Connect the requirement to a fundamental conservation law.

What is the difference between a coefficient and a subscript when balancing an equation, and why can you adjust one but not the other?

What is a reliable order in which to balance the elements in a typical molecular equation?

Hint: Save free elements (H2, O2, single metals) for last.

What are the five major types of chemical reactions?

What is the general form of a synthesis (combination) reaction?

What is the general form of a decomposition reaction?

What is a single-replacement reaction and what determines whether it will occur?

Hint: A ranking of element reactivity is involved.

What is the activity series and how is it used?

What is a double-replacement reaction and what driving forces make it proceed?

What products form when a hydrocarbon undergoes complete combustion?

How do you predict the products when a metal reacts with an acid?

What common decomposition patterns should you memorize?

What is a systematic strategy for classifying and predicting the products of an unknown reaction?

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Related Content

Study Topics

Reactions in SolutionRedox & BalancingStoichiometry

Elements

ZnCuMgFeNa

Reference Tables

Activity Series

Content Sources

Concept sections adapted from open educational resources under Creative Commons licensing:

  • OpenStax Chemistry 2e, Ch 4.1: Writing and Balancing Chemical Equations (CC BY 4.0)
  • OpenStax Chemistry 2e, Ch 4.2: Classifying Chemical Reactions (CC BY 4.0)
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