Types of Chemical Reactions

Addition and Decomposition Reactions

Addition reactions involve two or more reactants combining to form products. In decomposition reactions, a reactant decomposes into two or more products.

When two or more reactants combine to form a product, this is known as an addition reaction or synthesis reaction. The general equation for this reaction is:

\rm A+\rm B\rightarrow\rm{AB}

For example, when magnesium ribbon is burned in air, magnesium (Mg) reacts with oxygen gas (O₂) in air to form magnesium oxide.

{2{\rm{Mg}}}{(s)}+{\rm O}_2{(g)}\rightarrow2{\rm{MgO}}{(s)}

Note that this reaction is also a combustion reaction, a reaction in which a substance is burned in the presence of an oxidizer, and an oxidation-reduction reaction, a reaction in which oxidation states of two or more atoms change. The same reaction can be classified as multiple types. Another example is the reaction between nitrogen gas (N₂) and hydrogen gas (H₂), forming ammonia gas (NH₃).

{\rm N}_2{(g)}+3{\rm H}_2{(g)}\rightarrow2{\rm{NH}}_3{(g)}

This reaction, forming ammonia gas from atmospheric nitrogen, occurs only under very specific conditions. Ammonia is a nitrogen-based compound that is used to produce many important nitrogen-containing chemicals, such as artificial fertilizers.
A reaction in which a single compound breaks apart into two or more substances is called a decomposition reaction. The general equation for decomposition reaction is:

\rm{AB}\rightarrow\rm A+\rm B

There are three types of decomposition reactions. When a substance decomposes with application of heat, it is called thermal decomposition. For example, calcium carbonate (CaCO₃), when heated, decomposes into calcium oxide (CaO) and carbon dioxide (CO₂).

{\rm{CaCO}}_3{(s)}+\rm{heat}\rightarrow{\rm{CaO}}{(s)}+{\rm{CO}}_2{(g)}

When a substance decomposes because electricity passes through it, electrolytic decomposition occurs. Water (H₂O) decomposes into oxygen and hydrogen when electric current passes through it.

2{\rm H}_2{\rm O}{(l)}+\rm{electricity}\rightarrow2{\rm H}_2{(g)}+{\rm O}_2{(g)}

A substance can also decompose when exposed to light by a process called photodecomposition. Silver chloride (AgCl), for example, decomposes into silver (Ag) and chlorine gas (Cl₂) on exposure to sunlight.

2{\rm{AgCl}}{(s)}\rightarrow2{\rm{Ag}}{(s)}+{\rm{Cl}}_2{(g)}

Redox Reactions

In oxidation-reduction reactions, the oxidation states of two or more atoms change.

An oxidation state, also called an oxidation number, is a hypothetical charge assigned to an atom, ion, or polyatomic ion, indicating how many electrons have been lost (or gained). Chemical reactions all involve electrons. However, in a lot of chemical reactions, the oxidation states of atoms do not change. Chemical reactions in which an oxidation state of one or more atoms change are called oxidation-reduction reactions or redox reactions.

Oxidation-reduction reactions form the basis of electrochemistry, the branch of chemistry that studies batteries and other power cells. Combustion, burning a substance (often called fuel) with an oxidizer (often oxygen), is also a type of oxidation-reduction reaction. Consider the burning of methane (CH₄) with oxygen gas (O₂) to produce carbon dioxide (CO₂) and water (H₂O).

{\rm{CH}_4}({g})+{\rm{O}_2}({g})\rightarrow{\rm{CO}_2}({g})+2{\rm{H}_2\rm{O}}(l)

The oxidation state of each atom in an oxidation-reduction reaction is determined by its position in the periodic table and its bonding behavior in the reaction.

Oxidation States in ${\rm{CH}_4}({g})+{\rm{O}_2}({g})\rightarrow{\rm{CO}_2}({g})+2{\rm{H}_2\rm{O}}({l})$

CH₄	Hydrogen atoms are +1 when bonded to nonmetals. Carbon is –4 because it is coupled with four hydrogen atoms.
O₂	Since this is an elemental molecule, the oxidation state is zero.
CO₂	Oxygen in most compounds has an oxidation state of –2. Carbon is +4 because it is coupled with two oxygen atoms.
H₂O	Hydrogen atoms are +1 when bonded to nonmetals. Oxygen in most compounds is –2, which works out with two hydrogen atoms.

Overall, carbon's oxidation state went from –4 to +4. Hydrogen's oxidation state did not change and stayed +1. Oxygen's oxidation went from zero to –2.

Oxidation is a reaction that involves the removal of an electron from an atom. The carbon in methane in the example is oxidized. The opposite of oxidation is reduction. Reduction is a reaction that involves the addition of an electron to an atom. Oxygen gas in the example is reduced.

An oxidizing agent causes oxidation of another substance by stripping electrons from it, and it is reduced in the process. In the example, oxygen gas is the oxidizing agent. A reducing agent causes the reduction of another substance by giving it electrons, and it is oxidized in the process. In the example, carbon in methane is the reducing agent.

Oxidation always occurs with reduction. When an atom is oxidized, another must be reduced. The name oxidation comes from oxygen because oxygen is a common and strong oxidizing agent. Oxygen typically causes oxidation in other compounds and gets reduced in turn. There are other oxidizing agents, and some are stronger than oxygen. Fluorine, for example, is the most electronegative element and can cause oxidation in oxygen.

Oxidation-reduction reactions, such as combustion, can release energy as heat. It is possible to set up an oxidation-reduction reaction so that it releases energy as electricity. An oxidation-reduction reaction involves a transfer of electrons. If the reaction is separated into two parts, the electron transfer can occur over a wire. A flow of electrons causes electricity. Setting up an oxidation-reduction reaction this way involves setting it up as two half-reactions. A half-reaction is either the oxidation or the reduction part of an oxidation-reduction reaction.

Consider the reaction between magnesium (Mg) and copper oxide (CuO):

\rm{Mg}+\rm{CuO}\rightarrow\rm{Cu}+\rm{MgO}

The oxidation states of each element in the reaction are written above each atom in the chemical formulas.

\overset0{\rm M}\rm g+\overset{+2\;-2}{\rm{CuO}}\rightarrow\overset0{\rm Cu} +\overset{+2\;\;\;-2}{\rm{MgO}}

Notice how the oxidation states change. The oxidation state of magnesium changes from zero to +2. Because magnesium loses electrons, it is oxidized. The oxidation state of copper changes from +2 to zero. Because copper gains electrons, it is reduced. The oxidation state of oxygen remains as –2. Oxygen is not oxidized or reduced in this reaction.
It is possible to write the two half-reactions as:

\rm{Mg}\rightarrow\rm{Mg}^{2+}+2\rm e^-\;\;\;\text{(oxidation reaction)}

2\rm e^-+\rm{Cu}^{2+}\rightarrow\rm{Cu}\;\;\;\text{(reduction reaction)}

The electrons in this reaction move from magnesium to copper. Magnesium loses electrons, and copper gains them.

Neutralization (Acid-Base) Reactions

In neutralization reactions, acids and bases react with one another, forming salt and water.

The Arrhenius definition of acids and bases, proposed by Swedish chemist Svante Arrhenius, characterizes acids and bases depending on the ions produced by each in aqueous solution. An acid is a substance that produces hydrogen (H⁺) in solution, whereas a base produces hydroxide (OH^–) ions in solution.

A neutralization reaction is a reaction between an acid and a base. Acid-base reactions produce a salt and water.

The general equation is:

\text{Acid}+\text{Base}\rightarrow\text{Salt}+\text{Water}

Controlling acidity is important in living organisms. Neutralization reactions happen frequently in nature. Neutralization reactions are important in industry and agriculture as well.
For example, hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH), forming sodium chloride (NaCl) and water (H₂O).

{\rm {HCl}}{({aq})}+{\rm{NaOH}}{({aq})}\rightarrow{\rm{NaCl}}{({aq})}+{\rm H}_2{\rm O}{({l})}

In the example, hydrochloric acid is the acid. Sodium hydroxide is the base. The salt that forms is sodium chloride, or table salt.
Another example is the reaction of sulfuric acid (H₂SO₄) and potassium hydroxide (KOH), forming potassium sulfate (K₂SO₄) and water.

{\rm H}_2{\rm{SO}}_4{({aq})}+{2\rm{KOH}}{({aq})}\rightarrow{\rm K}_2{\rm{SO}}_4{({aq})}+2{\rm H}_2{\rm O}{(l)}

The acid in this example is sulfuric acid, the base is potassium hydroxide, and the salt is potassium sulfate.
A third example is the reaction of nitric acid (HNO₃) with sodium hydroxide, forming sodium nitrate (NaNO₃) and water.

{\rm {HNO}}_3{({aq})}+{\rm{NaOH}}{({aq})}\rightarrow{\rm {NaNO}}_3{({s})}+{\rm H}_2{\rm O}{(l)}

Single- and Double-Displacement Reactions

In single-displacement reactions, an atom or a group of atoms from a reactant is replaced with another atom or group of atoms in the products. In double-displacement reactions, two reactants exchange atoms or groups.

When an atom or a group of atoms from a reactant is replaced with another atom or group of atoms in the products, such a reaction is called a single-displacement reaction. The general form of a single-displacement reaction is:

\rm{AB}\;+\;\rm C\;\rightarrow\;\rm{AC}\;+\;\rm B

In a single-displacement reaction, C is more reactive than B. Otherwise, the reaction will not move forward. When zinc (Zn) reacts with hydrochloric acid (HCl), hydrogen (H) is replaced by zinc, forming zinc chloride (ZnCl₂).

{\rm{Zn}}{(s)}+2{\rm{HCl}}{(l)}\rightarrow{\rm{ZnCl}}_2{(s)}+{\rm H}_2{(g)}

In this reaction, the zinc replaces the hydrogen atom.
Another example is when hydrogen iodide (HI) reacts with chlorine gas (Cl₂), forming hydrogen chloride (HCl) and iodine (I₂)

2{\rm{HI}}{(g)}+{\rm{Cl}}_2{(g)}\rightarrow2{\rm{HCl}}{(g)}+{\rm I}_2{(s)}

In the reaction, chlorine gas replaces iodine.
When two reactants exchange atoms or groups, such a reaction is called a double-displacement reaction. These reactions usually take place in an aqueous state. The general form of this reaction is:

\rm{AB}+\rm{CD}\rightarrow\rm{AD}+\rm{BC}

Consider the reaction between sodium sulfide (Na₂S) and hydrochloric acid. Sodium chloride (NaCl), which is table salt, and hydrogen sulfide (H₂S) are produced by this reaction.

{\rm{Na}}_2{\rm S}({aq})+2{\rm{HCl}}({aq})\rightarrow2{\rm{NaCl}}({aq})+{\rm H}_2{\rm S}({g})

Sodium chloride is a soluble salt in this example. Single- and double-displacement reactions can produce a wide variety of salts. Some salts are not soluble in water and will precipitate. A precipitate is an insoluble product that settles as a residue at the bottom of the reaction vessel. A reaction that forms an insoluble salt, which forms as a solid in the reaction container, is called a precipitation reaction.

Solubility and Precipitate Formation

Double-displacement reactions do not always result in precipitates. Precipitation reactions are closely connected to solubility, the maximum amount of a substance that can be dissolved in another substance at specific conditions. A substance or dissolved material in a solution is called a solute. A substance that dissolves a material to form a solution is called a solvent. Water is a very common solvent. Solubility depends on the solute, on the solvent, and on conditions such as pressure and temperature.

A substance that has high solubility in a solvent under specific conditions is said to be soluble. Some substances are insoluble, which means they are incapable of being dissolved in another substance.

For example, when silver nitrate (AgNO₃) and sodium iodide (NaI) react with each other in aqueous forms, insoluble silver iodide (AgI) is formed as a precipitate.

{\rm A\rm g\rm N{\rm O}_3}{(aq)}+{\rm N\rm a\rm I}{(aq)}\rightarrow{\rm A\rm g\rm I}{(s)}+{\rm N\rm a\rm N{\rm O}_3}{(aq)}

This reaction can be written in ionic form as:

{\rm{Ag}^+}{({aq})}+{\rm{NO}_3}^-{({aq})}+{\rm{Na}^+}{(aq)}+{\rm I^-}{(aq)}\rightarrow{\rm{AgI}}{(s)}+{\rm{Na}^+}{(aq)}+{\rm{NO}_3}^-{(aq)}

When spectator ions (which are present in both the reactants and the products) are removed, the net ionic equation becomes:

{\rm{Ag}^+}{({aq})}+{\rm I^-}{({aq})}\rightarrow{\rm{AgI}}{(s)}

Similarly, when the solution of barium chloride (BaCl₂) reacts with sodium sulfate (Na₂SO₄), an insoluble precipitate of barium sulfate (BaSO₄) is formed.

{\rm{BaCl}}_2{({aq})}+{\rm{Na}}_2{\rm{SO}}_4{({aq})}\rightarrow{\rm{BaSO}}_4{(s)}+2{\rm{NaCl}}{({aq})}

This reaction can be written in ionic form as:

{\rm{Ba}^{2+}}({aq})+{\rm{SO}_4}^{2-}({aq})\rightarrow{\rm{BaSO}_4}({s})

The solid insoluble precipitate of barium sulfate is formed at the bottom of the test tube.

Precipitation Reaction

Magnesium reacts with copper sulfate to form copper as a precipitate.

Solubility Rules

Chemists have come up with a set of rules to determine if a salt is soluble in water or not. To predict if a precipitate will form, consider the solubility rules for salts in water. If two rules appear to contradict each other, the rule with the lower number takes precedence.

1. Group 1 element salts are soluble.

2. Salts with nitrate ions are soluble.

3. Salts of chloride, bromide, and iodide are soluble except for silver, lead, and mercury halides.

4. All silver salts are insoluble, except AgNO₃ and a few other rare exceptions.

5. Sulfate salts are soluble except for calcium, barium, lead, strontium, and silver salts.

6. Hydroxide salts of Group 1 elements are soluble. Hydroxide salts of Group 2 elements (calcium, strontium, and barium) are partially soluble. Hydroxide salts of transition metals are insoluble.

7. Sulfides of transition metals are insoluble.

8. Carbonates, phosphates, fluorides, and chromates are insoluble.

There are exceptions to these rules. However, one can predict precipitate formation based on the rules with fair success.

Multiple-Classifications Reactions

A reaction may belong to several different categories.

Reactions can be classified into different types, but often a reaction can belong to more than one category. For example, almost all acid-base reactions are also double-displacement reactions. Many redox reactions are also combination reactions. Consider the reaction in which nitric acid (HNO₃) reacts with sodium hydroxide (NaOH), a base, and forms a salt, sodium nitrate (NaNO₃).

{\rm{HNO}}_3{({aq})}+{\rm{NaOH}}\,{({aq})}\rightarrow{\rm{NaNO}}_3{({s})}+{\rm H}_2{\rm O}{({l})}

This is an acid-base reaction and a double-displacement reaction at the same time.
Another example is the reaction of calcium (Ca) with fluorine gas (F₂) to form calcium fluoride (CaF₂).

{\rm{Ca}}({s})+{\rm F}_2({g})\rightarrow{\rm{CaF}}_2({s})

This is a combination, as well as an oxidation-reduction reaction, because calcium's oxidation state has changed from zero to +2 and fluorine's oxidation state has changed from zero to –1.
Zinc (Zn) reacts with hydrochloric acid (HCl) to form zinc chloride (ZnCl₂) and hydrogen gas (H₂).

{\rm{Zn}}({s})+2{\rm{HCl}}({aq})\rightarrow{{\rm{ZnCl}}_2}({aq})+{\rm H}_2({g})

This is a single-displacement reaction, as well as an oxidation-reduction reaction, because zinc's oxidation state has changed from zero to +2 and hydrogen's oxidation state has changed from +1 to zero.
Sodium hydroxide (NaOH), a strong base, reacts with hydrochloric acid (HCl) and forms sodium chloride (NaCl) and water (H₂O).

{\rm{NaOH}}({aq})+{\rm{HCl}}({aq})\rightarrow{\rm{NaCl}}({aq})+{\rm {H}_2}{\rm O}({l})

This is a neutralization reaction as well as a double-displacement reaction.

<Ionic Equations >Suggested Reading

Reactions in Chemistry

Contents

Types of Chemical Reactions

Addition and Decomposition Reactions

Redox Reactions

Oxidation States in ${\rm{CH}_4}({g})+{\rm{O}_2}({g})\rightarrow{\rm{CO}_2}({g})+2{\rm{H}_2\rm{O}}({l})$

Neutralization (Acid-Base) Reactions

Single- and Double-Displacement Reactions

Solubility and Precipitate Formation

Precipitation Reaction

Solubility Rules

Multiple-Classifications Reactions

Reactions in Chemistry

Contents

Types of Chemical Reactions

Addition and Decomposition Reactions

Redox Reactions

Oxidation States in CH4(g)+O2(g)→CO2(g)+2H2O(l){\rm{CH}_4}({g})+{\rm{O}_2}({g})\rightarrow{\rm{CO}_2}({g})+2{\rm{H}_2\rm{O}}({l})CH4​(g)+O2​(g)→CO2​(g)+2H2​O(l)

Neutralization (Acid-Base) Reactions

Single- and Double-Displacement Reactions

Solubility and Precipitate Formation

Precipitation Reaction

Solubility Rules

Multiple-Classifications Reactions

Oxidation States in ${\rm{CH}_4}({g})+{\rm{O}_2}({g})\rightarrow{\rm{CO}_2}({g})+2{\rm{H}_2\rm{O}}({l})$