Chemical Reactions

Chemistry is boring without reactions; everything just sits there. There are only 50 or so elements of which we routinely make use, but a practically infinite number of ways in which to combine those to make new materials with different, sometimes unexpected properties. We're still finding ways to put elements and compounds together (synthesis) and to take them apart (decomposition). Those processes necessarily involve reactions.

This section will review some elementary reaction types, then we'll talk about some more complex kinds. Many of the reactions written below contain numerical coefficients in front of molecules, like 2H2O, which simply means "two molecules of water. Don't worry too much about those for now, you'll learn more when you learn to balance chemical equations.

1. Synthesis reactions

The reaction in the green box below is a prototype (that means A's and B's stand for elements) of a typical synthesis reaction, a reaction in which atoms or smaller molecules combine to form one or more larger or more complicated compound(s).

Before we go on to some examples, let's use this reaction to take care of some basic reaction terminology. The compound(s) on the left are called reactant(s). Those on the left are product(s). We separate different sets of reactants and products with + signs.

About arrows →

We generally only use a single arrow ( ) to indicate a few reactions where the reverse reaction is very improbable, like the dissociation of a strong acid.

More often we use a double arrow ( ⇌ ) to show that the reaction actually proceeds in both directions at the same time. There will be more to say about that in the section on equilibrium, but for now, you should start using the double arrow to write most reactions.

synthesis reaction

Examples of synthesis reactions

Here are a few examples of synthesis reactions. Compare them to the model reaction above to make sure you get the idea. For each reaction there's a version written in an English sentence on the right, so that you might be able to pick up on the meaning more easily:

Reaction What it means*
Fe + O2 ⇌ FeO2 Iron combines with oxygen to form iron oxide
H2 + Br2 ⇌ 2 HBr Diatomic hydrogen combines with diatomic bromine to form two molecules of hydrogen bromide
2 NH3 + H2O + CO2 ⇌ (NH4)2CO3 Two molecules of ammonia, a water and a carbon dioxide combine to form a molecule of ammonium carbonate.
CaO + SO2 ⇌ CaSO3 Calcium oxide and sulfur dioxide react to form calcium sulfite.

A note about the sentences above (and below): These aren't the only ways to describe these reactions, just ones I though appropriate at the moment. You might do a better job.

Synthesis reactions

In a synthesis reaction, two or more simpler components are joined to form a more complex compound.

Solids, liquids & gases

Often we write the state (solid, liquid, gas, aqueous) of a compound in a reaction using parenthesis and the letters (s), (l), (g) & (aq). Aqueous solutions are solutions in which water is the solvent. They are so ubiquitous in our chemistry that they get the special designation (aq). Here's the table above, rewritten with the added state information:

Fe (s) + O2 (g) ⇌ FeO2 (s) solid Iron combines with oxygen gas to form solid iron oxide
H2 (g) + Br2 (g) ⇌ 2 HBr (g) Hydrogen gas combines with bromine gas to form two molecules of gaseous hydrogen bromide
2 NH3 (g) + H2O (l) + CO2 (g) ⇌ (NH4)2CO3 (aq) Two molecules of gaseous ammonia, a liquid water molecule and a molecule of carbon dioxide gas combine to form a molecule of aqueous ammonium carbonate.

Aqueqous (ā' · kwee · us) solutions are mixtures of soluble ionic or nonionic compounds and water. They are pure substances.

2. Decomposition reactions

Decomposition reactions are really just the opposite of synthesis. Something more complicated "falls apart" into less complicated things. Decomposition is usually what we're talking about when we speak of "shelf life", like for drugs or some foods.

Decomposition that is catalyzed by visible or ultraviolet light is why many drugs come in brown plastic bottles (the brown color blocks a lot of visible light and plastics absorb most UV).

Examples of decomposition reactions

AgO2 (aq) ⇌ Ag(s) + O2(g) Aqueous silver oxide decomposes into solid silver metal and oxygen gas.
H2SO4(aq) ⇌ 2H+(aq) + SO42-(aq) Sulfuric acid dissociates into two protons plus a sulfate ion.
H2O2 (aq) ⇌ H2(g) + O2(g) Aqueous hydrogen peroxide decomposes into hydrogen and oxygen gas.

3. Single-displacement reactions

Displacement reactions might also be called substitution reactions. In a displacement reaction, one chemical moiety from one reactant takes the place of another in the other reactant.

You can think of single-displacement reactions as ones in which a molecule donates part of itself to another atom or molecule. Here's the picture →



moy'   ·   it   ·   ee

The word moiety is sometimes used in chemistry to identify one reactant or product.

Literally, it means each of the two parts into which a thing can be divided.

Examples of single-displacement reactions

Cu (s) + 2AgNO3 (aq) ⇌ CuNO3(aq) + 2Ag(s) When solid copper and aqueous silver nitrate react, copper replaces silver in the nitrate compound, with solid silver remaining as a product.
Fe(s) + 2HCl (g) ⇌ FeCl2(aq) + H2(g) When solid iron is treated with hydrochloric acid, two chlorines from the acid combine with iron to form iron (II) chloride and hydrogen gas.
2Al (s) + Fe2O3 (aq) ⇌ Al2O3(aq) + 2Fe(s) Two aluminum atoms react with one atom of Iron (III) oxide in aqueous solution to form aluminum (III) oxide and solid iron.

4. Double-displacement reactions

The double displacement reaction is a pretty obvious extension of single displacement. It requires two binary compounds, each of which exchanges one of its parts with the other. Once you understand double displacement, it's possible to imagine all kinds of displacement reactions in which more complicated compounds swap one or more parts. Here is the double displacement model:

Examples of double-displacement reactions

Pb(NO3)2 (aq)+ 2NaCl (aq) ⇌ 2NaNO3(aq)+ PbCl2(s) Lead (II) nitrate reacts with the salt sodium chloride to produce sodium nitrate and lead (II) chloride.
H2SO4(aq) + 2LiOH (aq) ⇌ Li2SO4(aq) + 2H2O(l) Aqueous sulfuric acid reacts with lithium hydroxide to produce lithium sulfate and water.
AgNO3 (aq) + HCl (g) ⇌ AgCl(s) + HNO3(aq) Aqueous silver nitrate reacts with hydrochloric acid to produce a solid precipitate of silver chloride and aqueous nitric acid.

5. Acid-Base neutralization

Acids and bases are covered in another section, so this kind of reaction might seem a little fuzzy. You can come back to it later. For now, consider an acid to be an ionic compound that yields H+ ions and a base to be an ionic compound that yields OH- ions. Excess H+ ions are what make a solution acidic and excess OH- ions are what makes it basic. OH- and H+ can react to form H2O, which is (exactly) neither acidic nor basic—water is, by definition, neutral.

The products of neutralization are always a salt (a non-acidic, non-basic ionic compound) and water. Here's what it looks like:

"A" is an abbreviation for the anion of an acid, such as Cl- or SO42-, and "B" is the usual abbreviation for the cation of a base, such as Na+

Examples of acid-base neutralization reactions

HCl (aq) + NaOH (aq) ⇌ NaCl(aq) + H2O(l) Hydrochloric acid is neutralized with an equimolar amount of sodium hydroxide in aqueous solution, forming the salt NaCl and liquid water at pH 7.
H3PO4(aq) + 3KOH (aq) ⇌ K3PO4(aq) + 3H2O(l) Phosphoric acid is completely neutralized with potassium hydroxide to form potassium phosphate and liquid water.
2HCOOH (aq) + Ba(OH)2 (aq)
2HCOOBa2(aq) + 2H2O(l)
Formic acid (HCOOH) is neutralized with barium hydroxide to yield barium formate (a salt) and water.

Writing neutralization reactions — identifying the salt

Here's a trick for identifying the products in acid-base neutralization reactions. Let's say you want to write a reaction for the neutralization of perchloric acid, HClO4, and calcium hydroxide, Ca(OH)2.

First write dissociation reactions for the acid and base. These are really just decomposition reactions; salts like these tend to come apart into separate ions when placed in water. Here are the reactions:


HClO4   →   H+ + ClO4-


Ca(OH)2   →   Ca2+ + 2 OH-

Now we know that one of the products is H2O. that's obviously going to be formed from the H+ and OH- ions, so the salt is just a combination of the remaining ions, Ca2+ and ClO4-, in the right proportions to make a neutral compound,


Having a list of common molecular (polyatomic) ions around will help, too.

6. Combustion reactions

Combustion is rapid oxidation or burning. It is the (usually) rapid combination of oxygen with a hydrocarbon (composed only of C and H) or an oxy-hydrocarbon (also contains oxygen). Combustion generally releases a great deal of the energy stored in the chemical bonds of a molecule.

The model reaction below is a general form for balancing any hydrocarbon reaction. It's interesting, but I wouldn't spend any time trying to memorize it. It's better just to learn to balance any kind of reaction quickly and be done with it.

Examples of combustion reactions

CH4 (g) + 2O2 (g) ⇌ CO2(g) + 2H2O(g) Methane (CH4) gas is combusted in the presence of oxygen to yield carbon dioxide and water.
2C2H5OC2H5(l) + 13O2 (g) ⇌ 8CO2(g) + 10H2O(g) Diethyl ether is burned in the presence of oxygen to yield carbon dioxide and water.
2C6H6 (g) + 15O2 (g) ⇌ 12CO2(g) + 6H2O(g) Two molecules of benzene (C6H6) are burned in the presence of excess oxygen to yield 12 molecules of CO2 and six molecules of H2O.

Dimethyl ether, burned in the second reaction above, is an oxy-hydrocarbon. Think of it as a hydrocarbon that carries around a bit of its own oxygen for combustion. In fact, molecules like these can help to ensure that combustion reactions go to completion, especially in unfavorable conditions — like cold weather. Diethyl ether is added to gasoline in some states in winter to help the gas combust fully, avoiding other products that are more harmful to life, such as carbon monoxide (CO).

7. Rearrangement or isomerization reactions

Some molecules can undergo an internal rearrangement of their structure under certain circumstances. Molecules that have the same number and type of atoms, but different bonding arrangem#ents are called isomers, and interconversion between isomers is called isomerization. Here's an example, the rearrangement of n-butane to isobutane. This diagram is actually a shorthand; I've omitted all of the hydrogens that would be bound to the carbons – each would have to have four bonds.

Here's another example, the rearrangement of a 6-carbon sugar molecule between two forms. →

Practice problems

Determine they type of each reaction below. Roll over or tap the equation for the answer.


2 C6H6 + 15 O2 ⇌ 12 CO2 + 6 H2O


Combustion: Combines a hydrocarbon with oxygen to produce CO2 and water.


NaBr + H3PO4 ⇌ Na3PO4 + HBr


Double displacement: The ions are H+, PO43-, Na+, and Br-.


Fe3+ + 3 NaBr ⇌ FeBr3 + 3 Na+


Single displacement: The ions are Fe3+, Na+, and Br-. The iron ion is substituted for the sodium ion, and the number of bromines has to change because of the difference in cation charge.


Mn + O2 ⇌ MnO2


Synthesis: Two compounds are joined into one.


Na2CO3 ⇌ Na2O + CO2


Decomposition: One compound breaks into two or more separate compounds.


3 Ca(OH)2 + Al2(SO4)3 ⇌ 3 CaSO4 + 2 Al(OH)3


Double displacement: The ions are Ca2+, OH-, Al3+, and the molecular ion SO42-.


2 PbSO4 ⇌ 2 PbSO3 + O2


Decomposition: The compound on the left breaks into two by losing an oxygen molecule (O2).


H2SO4 + 2 NH4OH ⇌ 2 H2O + (NH4)2SO4


Double displacement: The ions are H+, SO42-, OH-, and NH4+.


CaSO4 + Mg(OH)2 ⇌ Ca(OH)2 + MgSO4


Double displacement: The ions are Ca2+, SO4, Mg2+, and OH-.


C2H4 + O2 ⇌ CO2 + H2O


Combustion: Combines a hydrocarbon with oxygen to produce CO2 and water.

Oxidation-reduction (redox)

The term "redox" is commonly used to describe oxidation-reduction reactions. Redox reactions are ones in which electrons are transferred from one reactant to another. The reactant that loses electrons is said to be oxidized and the one that gains them is said to be reduced.

I know that sounds a bit like opposite world, but it's a historical artifact that we're stuck with: gaining electrons = "reduction."

We usually write redox reactions in terms of half reactions, representing the oxidation half reaction and the reduction half reaction separately. They can then be added to get the full redox reaction.

For example, the zinc-copper redox reaction looks like this:

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

The oxidation half reaction, in which solid zinc loses two electrons, is

Zn(s)   →   Zn2+(aq) + 2e-

Notice that in the half reactions we incorporate electrons as a reactant. The reduction half reaction, in which copper ions gain two electrons to convert to solid copper, is

Cu2+(aq) + 2e-   →   Cu(s)

Two electrons are transferred in each reaction. When the two half reactions are added, there are two electrons on either side of the reaction arrow, and we simply cancel them algebraically.

The sections on electrochemistry and redox balancing will give you a much better look at redox reactions; I present them here just for completeness of this section.

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