3.1 Chemical Reactions & Chemical Equations
PRE-HEALTH PREP
Chemical Reactions and Equations
General Chemistry 3.1
Lesson Overview
Chemical reactions in general chemistry fall into five fundamental categories: combination, decomposition, combustion, double replacement, and single replacement. Each reaction type follows a distinct pattern that determines how products form — from ionic compounds assembling by charge balance in combination reactions to hydrocarbons yielding carbon dioxide and water during combustion. Product prediction strategies and equation balancing techniques differ by reaction type, with key principles including writing formulas from ionic charges rather than copying reactant subscripts, and saving oxygen for last when balancing combustion equations.
Table of Contents
Introduction to Chemical Reactions and Equations
This lesson opens the chapter on stoichiometry — one of the most important (and, admittedly, most anxiety-inducing) topics in general chemistry. Stoichiometry revolves around the mole and Avogadro's number, but before we can dive into those calculations, we need to establish a solid foundation with chemical reactions and equations. Consider this lesson the on-ramp. By the time you reach the end, you will either genuinely enjoy this material or, at minimum, find it far less daunting than you expected.
The goals for this lesson are straightforward: learn to recognize 5 types of chemical reactions, predict products for four of them, and balance the resulting equations.
The 5 Types of Chemical Reactions
Chemistry organizes an enormous variety of reactions into recognizable patterns. The five reaction types you must know for this chapter are:
- Combination reactions — two or more reactants merge into a single product
- Decomposition reactions — a single reactant breaks apart into multiple products
- Combustion reactions — a substance reacts with oxygen gas, releasing energy
- Double replacement (double displacement) reactions — cations and anions swap partners between two ionic compounds
- Single replacement (single displacement) reactions — one element displaces another element from a compound
Pattern recognition tip: Before you do anything else with a reaction, identify which type it is. The type tells you what products to expect and how to approach balancing. Recognizing the pattern is half the work.
Combination Reactions
A combination reaction is exactly what its name implies: two or more reactants combine to form a single product. The hallmark is simple — multiple reactants in, one product out.
The most common scenario is combining an elemental metal with an elemental nonmetal. When these two species meet, they form an ionic compound. Consider sodium metal reacting with chlorine gas — a classic combination reaction.
Writing the Correct Product Formula
A critical mistake students make is to simply copy the elemental formulas onto the product side without thinking about ionic charges. An ionic compound must have zero net charge, so you need to figure out how many of each ion are required before writing anything down.
Recall how to predict ionic charges from position on the periodic table:
- Sodium (Na) sits one position beyond the previous noble gas, so it loses one electron and carries a +1 charge as an ion.
- Chlorine (Cl) sits one position short of the next noble gas, so it gains one electron and carries a −1 charge as an ion.
With equal and opposite charges, you need one Na⁺ for every one Cl⁻, giving the formula NaCl — not NaCl₂, and certainly not Na₂Cl.
Common mistake — carrying subscripts across the arrow: Students often write the product formula by importing subscripts from the reactant side. Do not do this. The subscripts in the product are determined solely by the need for the ionic charges to balance to zero. Determine the formula from charges first, then write it.
Balancing the Combination Reaction
Once you have the correct product formula, you balance the equation using coefficients — never by changing subscripts in a formula, because that would alter the compound itself.
Example: Sodium + Chlorine
Start: Na + Cl₂ → NaCl
The sodium atoms are balanced (one on each side), but chlorine is not: two Cl atoms on the left from Cl₂, only one on the right. Place a coefficient of 2 in front of NaCl:
Na + Cl₂ → 2 NaCl
Now sodium is unbalanced — two on the right, one on the left. Place a coefficient of 2 in front of Na:
2 Na + Cl₂ → 2 NaCl
Check: 2 Na on each side, 2 Cl on each side. Balanced. This is a combination reaction because two reactants (Na and Cl₂) yield one product (NaCl).
Decomposition Reactions
A decomposition reaction is the conceptual inverse of a combination reaction. Instead of multiple reactants merging into one product, a single reactant breaks apart into two or more products. The identifying signature: one reactant in, multiple products out.
Most decomposition reactions are thermal decompositions, meaning they require heat to proceed. In chemical notation, heat is represented by the Greek letter delta (Δ) written above the reaction arrow.
Example: Decomposition of Calcium Carbonate
CaCO₃ →(Δ) CaO + CO₂
Calcium carbonate (CaCO₃) decomposes upon heating into calcium oxide (CaO) and carbon dioxide gas (CO₂).
Verify the balance: 1 Ca = 1 Ca; 1 C = 1 C; 3 O on the left, 1 + 2 = 3 O on the right. The equation is balanced as written — no coefficients needed. (When no coefficient appears in front of a species, a coefficient of 1 is implied.)
Practice Problem: Balance the following decomposition reaction.
__CaCO3(s) ➔ __CaO(s) + __CO2(g)
Fill in the correct coefficients (blanks indicated by __) to balance the equation. Hint: each element appears only once on each side — count atoms carefully before assigning coefficients.
Note on product prediction: Unlike combination, combustion, and replacement reactions, the products of a decomposition reaction are not always predictable from a simple rule. For now, products will be given to you. The important skill is recognizing the pattern — one reactant, multiple products — and verifying the balance.
Combustion Reactions
A combustion reaction is the reaction of a substance with oxygen gas (O₂). At its core, combustion is burning — the rapid release of energy as heat and light when a fuel reacts with oxygen.
Why can't you start a fire on the moon? There is no atmosphere, and therefore no oxygen. That same principle explains why covering a beaker with a watch glass extinguishes a small flame — once the available oxygen inside is consumed, the combustion stops. Fire blankets used in chemistry labs work on the same logic: wrap the flame and cut off its oxygen supply.
Combustion of Hydrocarbons: 2 Predictable Products
In general chemistry, you will almost always encounter the combustion of hydrocarbons — compounds composed entirely of carbon and hydrogen. The products are predictable every time:
- All carbon atoms in the fuel become carbon dioxide (CO₂).
- All hydrogen atoms become part of water (H₂O).
Hydrocarbon + O₂ → CO₂ + H₂O. Memorize this pattern — product prediction for hydrocarbon combustion is non-negotiable.
Balancing a Combustion Reaction: 3-Step Method
Combustion reactions can be among the trickiest to balance because oxygen appears in multiple products. The strategy: save oxygen for last. Balance carbon first, then hydrogen, and handle oxygen only after the other elements are locked in.
Example: Combustion of Propane (C₃H₈)
Start: C₃H₈ + O₂ → CO₂ + H₂O
Step 1 — Balance carbon. Propane has 3 carbons; each CO₂ contains 1. Place a 3 in front of CO₂:
C₃H₈ + O₂ → 3 CO₂ + H₂O
Step 2 — Balance hydrogen. Propane has 8 hydrogens; each H₂O contains 2. Place a 4 in front of H₂O:
C₃H₈ + O₂ → 3 CO₂ + 4 H₂O
Step 3 — Balance oxygen. On the right: 3 × 2 = 6 oxygens from CO₂, plus 4 × 1 = 4 oxygens from H₂O, for a total of 10. Place a 5 in front of O₂:
C₃H₈ + 5 O₂ → 3 CO₂ + 4 H₂O
Check: 3 C = 3 C; 8 H = 8 H; 10 O = 10 O. Balanced.
Practice Problem: Balance the combustion of propane.
__C3H8(g) + __O2(g) ➔ __CO2(g) + __H2O(g)
Fill in the correct coefficients (blanks indicated by __) using the 3-step method: balance carbon first, then hydrogen, then oxygen last.
Strategy — avoid starting with oxygen: Oxygen appears in both CO₂ and H₂O on the product side, which makes it impossible to fix independently. Always balance carbon and hydrogen first, then count up the total oxygen needed and set the O₂ coefficient accordingly.
Double Replacement (Double Displacement) Reactions
Double replacement reactions — also called double displacement, metathesis, or exchange reactions — occur when two ionic compounds in solution swap their cation-anion partners. Each cation ends up with the other compound's anion. These reactions typically take place in aqueous solution (water as solvent), which is why they are also classified as aqueous reactions.
You will revisit these reactions with much greater depth in Chapter 4, where you will add phase labels and apply solubility rules. For now, the focus is on recognizing the pattern and correctly predicting product formulas.
Predicting the Products
The rule is simple: cations and anions trade partners. Given silver nitrate (AgNO₃) and sodium sulfide (Na₂S), identify the ions:
- Silver nitrate: Ag⁺ (cation) + NO₃⁻ (anion)
- Sodium sulfide: Na⁺ (cation) + S²⁻ (anion)
Swap the anions: Ag⁺ pairs with S²⁻ to form silver sulfide; Na⁺ pairs with NO₃⁻ to form sodium nitrate.
Writing the Correct Product Formulas
As always, the formula of each ionic product must be electrically neutral.
For silver sulfide: Silver is one of the four special transition metals with a fixed charge — Ag is always +1. Sulfur as a monatomic anion gains two electrons (to match the argon configuration), giving S²⁻. One Ag⁺ (+1) cannot balance one S²⁻ (−2), so two silver ions are required: Ag₂S (total charge: +2 + −2 = 0).
For sodium nitrate: Na⁺ (+1) and NO₃⁻ (−1) balance in a 1:1 ratio. The formula is simply NaNO₃.
Polyatomic ions stay intact. The nitrate ion (NO₃⁻) moves as a complete unit from one compound to the other — you do not break it apart or separately balance individual nitrogen and oxygen atoms. Treat polyatomic ions as single indivisible species when writing formulas and when balancing.
Example: AgNO₃ + Na₂S → Ag₂S + NaNO₃ (unbalanced)
The product side shows 2 Ag in Ag₂S. To supply 2 Ag⁺ on the reactant side, place a 2 in front of AgNO₃ — this also delivers 2 NO₃⁻ groups. To consume those 2 nitrate groups on the product side, place a 2 in front of NaNO₃ — which also gives us 2 Na⁺. Check the reactant side: Na₂S already provides 2 Na⁺. Everything balances:
2 AgNO₃ + Na₂S → Ag₂S + 2 NaNO₃
Verify: 2 Ag = 2 Ag; 2 Na = 2 Na; 1 S = 1 S; 2 NO₃ = 2 NO₃. Balanced.
Looking ahead to Chapter 4: When you add phase labels, sodium nitrate (NaNO₃) will be marked aqueous (it dissolves readily), while silver sulfide (Ag₂S) is insoluble and precipitates as a solid. You will learn solubility rules in Chapter 4 to make these predictions yourself.
Single Replacement (Single Displacement) Reactions
A single replacement reaction occurs when a free element displaces an ion of the same type (metal displaces metal, or nonmetal displaces nonmetal) from a compound. Where double replacement involves two compounds exchanging partners, single replacement involves one lone element muscling its way in. These reactions also typically take place in aqueous solution and are, at their core, oxidation-reduction (redox) reactions — as is combustion. Redox chemistry will be covered thoroughly in Chapter 4; for now, focus on recognizing the pattern and predicting products.
Predicting the Products
Think of it this way: copper(II) chloride (CuCl₂) has copper and chloride bonded together as a pair. Elemental zinc arrives as a single, unattached species. Zinc can adopt a +2 ionic charge — the same role copper was playing — so it displaces copper from the compound and bonds with the chloride ions instead. Copper is released as a free metal.
Example: Zinc displaces copper from copper(II) chloride
Start: Zn + CuCl₂ → ? + ?
Zinc as an ion is always Zn²⁺ (a fixed-charge species). Chloride is Cl⁻. To balance Zn²⁺, two Cl⁻ are needed: the product is ZnCl₂. Copper is released as elemental copper metal, Cu(s).
Zn + CuCl₂ → ZnCl₂ + Cu
Check: 1 Zn = 1 Zn; 1 Cu = 1 Cu; 2 Cl = 2 Cl. Balanced as written — no coefficients needed.
Balancing single replacement reactions is generally more straightforward than balancing combustion reactions. Because only one element is swapping in and out, the stoichiometry rarely becomes complicated. Single replacement reactions will get more involved when you revisit redox in Chapter 4, but for now they are among the easiest to balance.
Key Takeaways
- 5 reaction types to know: combination, decomposition, combustion, double replacement, and single replacement. Recognizing which type you are looking at is the first step in predicting products and balancing equations.
- Combination reactions have two or more reactants producing a single product. When a metal combines with a nonmetal, an ionic compound forms — write the product formula by balancing ionic charges to zero, not by copying reactant subscripts.
- Decomposition reactions have one reactant breaking into multiple products. Most require heat (indicated by Δ over the arrow). Product prediction requires additional information; focus on recognizing the pattern and balancing.
- Combustion reactions are always a substance reacting with O₂. For hydrocarbons, the products are always CO₂ and H₂O. When balancing, save oxygen for last — balance carbon first, then hydrogen, then count total oxygen needed.
- Double replacement reactions swap cation-anion partners between two ionic compounds. Polyatomic ions (such as NO₃⁻) remain intact throughout. Use ionic charges to write correct product formulas, then balance with coefficients.
- Single replacement reactions have a lone element displacing an ion from a compound. Both combustion and single replacement are classified as redox reactions, though the full redox treatment comes in Chapter 4.
- Balancing rules that never change: only coefficients may be adjusted when balancing — subscripts within a formula are fixed. A missing coefficient always implies 1.