9.4: Classifying Chemical Reactions- Take One
- Page ID
- 451549
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The chemical reactions we have described are only a tiny sampling of the infinite number of chemical reactions possible. How do chemists cope with this overwhelming diversity? How do they predict which compounds will react with one another and what products will be formed? The key to success is to find useful ways to categorize reactions. Familiarity with a few basic types of reactions will help you to predict the products that form when certain kinds of compounds or elements come in contact.
Many chemical reactions can be classified into one or more of four basic types: combination, decomposition, single replacement, or double replacement. The general forms of these four kinds of reactions are summarized in Table \(\PageIndex{1}\), along with examples of each.
| Name of Reaction | General Form | Examples |
|---|---|---|
|
Single Replacement |
AB + C → AC + B | ZnCl2(aq)+ Mg(s) → MgCl2(aq)+ Zn(s) |
| Double Replacement | AB + CD → AD + CB | BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq) |
| Combination | A + B → AB |
CO2(g) + H2O(l) → H2CO3(aq) N2(g) + 2O2(g)→ 2NO2(g) |
| Decomposition | AB → A + B | CaCO3(s) → CaO(s) + CO2(g) |
The classification scheme is only for convenience; the same reaction can be classified in different ways, depending on which of its characteristics is most important. We will see another way of classifying reactions on the next page.
Combination Reactions
A combination reaction is a reaction in which two or more substances combine to form a single new substance. Combination reactions can also be called synthesis reactions. The general form of a combination reaction is:
\[\ce{A} + \ce{B} \rightarrow \ce{AB}\]
One combination reaction is two elements combining to form a compound. Solid sodium metal reacts with chlorine gas to produce solid sodium chloride.
\[2 \ce{Na} \left( s \right) + \ce{Cl_2} \left( g \right) \rightarrow 2 \ce{NaCl} \left( s \right) \nonumber \]
Notice that in order to write and balance the equation correctly, it is important to remember the seven elements that exist in nature as diatomic molecules (\(\ce{H_2}\), \(\ce{N_2}\), \(\ce{O_2}\), \(\ce{F_2}\), \(\ce{Cl_2}\), \(\ce{Br_2}\), and \(\ce{I_2}\)).
One type of combination reaction that occurs frequently is the reaction of an element with oxygen to form an oxide. Metals and nonmetals both react readily with oxygen under most conditions. Magnesium reacts rapidly and dramatically when ignited, combining with oxygen from the air to produce a fine powder of magnesium oxide:
\[2 \ce{Mg} \left( s \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{MgO} \left( s \right) \nonumber \]
Sulfur reacts with oxygen to form sulfur dioxide:
\[\ce{S} \left( s \right) + \ce{O_2} \left( g \right) \rightarrow \ce{SO_2} \left( g \right) \nonumber \]
When nonmetals react with one another, the product is a molecular compound. Often, the nonmetal reactants can combine in different ratios and produce different products. Sulfur can also combine with oxygen to form sulfur trioxide:
\[2 \ce{S} \left( s \right) + 3 \ce{O_2} \left( g \right) \rightarrow 2 \ce{SO_3} \left( g \right) \nonumber \]
Transition metals are capable of adopting multiple positive charges within their ionic compounds. Therefore, most transition metals are capable of forming different products in a combination reaction. Iron reacts with oxygen to form both iron (II) oxide and iron (III) oxide:
\[2 \ce{Fe} \left( s \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{FeO} \left( s \right) \nonumber \]
\[4 \ce{Fe} \left( s \right) + 3 \ce{O_2} \left( g \right) \rightarrow 2 \ce{Fe_2O_3} \left( s \right) \nonumber \]
Combination reactions can also take place when an element reacts with a compound to form a new compound composed of a larger number of atoms. Carbon monoxide reacts with oxygen to form carbon dioxide according to the equation:
\[2 \ce{CO} \left( g \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{CO_2} \left( g \right)\]
Two compounds may also react to form a more complex compound. A very common example is the reactions of oxides with water. Calcium oxide reacts readily with water to produce an aqueous solution of calcium hydroxide:
\[\ce{CaO} \left( s \right) + \ce{H_2O} \left( l \right) \rightarrow \ce{Ca(OH)_2} \left( aq \right)\]
Sulfur trioxide gas reacts with water to form sulfuric acid. This is an unfortunately common reaction that occurs in the atmosphere in some places where oxides of sulfur are present as pollutants. The acid formed in the reaction falls to the ground as acid rain.
\[\ce{SO_3} \left( g \right) + \ce{H_2O} \left( l \right) \rightarrow \ce{H_2SO_4} \left( aq \right)\]
Decomposition Reactions
A decomposition reaction is a reaction in which a compound breaks down into two or more simpler substances. The general form of a decomposition reaction is:
\[\ce{AB} \rightarrow \ce{A} + \ce{B}\]
Most decomposition reactions require an input of energy in the form of heat, light, or electricity.
Binary compounds are compounds composed of just two elements. The simplest kind of decomposition reaction is when a binary compound decomposes into its elements. Mercury (II) oxide, a red solid, decomposes when heated to produce mercury and oxygen gas:
\[2 \ce{HgO} \left( s \right) \rightarrow 2 \ce{Hg} \left( l \right) + \ce{O_2} \left( g \right)\]
Video \(\PageIndex{2}\): Mercury (II) oxide is a red solid. When it is heated, it decomposes into mercury metal and oxygen gas.
A reaction is also considered to be a decomposition reaction even when one or more of the products are still compounds. A metal carbonate decomposes into a metal oxide and carbon dioxide gas. For example, calcium carbonate decomposes into calcium oxide and carbon dioxide:
\[\ce{CaCO_3} \left( s \right) \rightarrow \ce{CaO} \left( s \right) + \ce{CO_2} \left( g \right)\]
Metal hydroxides decompose on heating to yield metal oxides and water. Sodium hydroxide decomposes to produce sodium oxide and water:
\[2 \ce{NaOH} \left( s \right) \rightarrow \ce{Na_2O} \left( s \right) + \ce{H_2O} \left( g \right)\]
Some unstable acids decompose to produce nonmetal oxides and water. Carbonic acid decomposes easily at room temperature into carbon dioxide and water:
\[\ce{H_2CO_3} \left( aq \right) \rightarrow \ce{CO_2} \left( g \right) + \ce{H_2O} \left( l \right)\]
Single Replacement Reactions
A third type of reaction is the single replacement reaction, in which one element replaces a similar element in a compound. The general form of a single-replacement (also called single-displacement) reaction is:
\[\ce{A} + \ce{BC} \rightarrow \ce{AC} + \ce{B}\]
In this general reaction, element \(\ce{A}\) is a metal and replaces element \(\ce{B}\), also a metal, in the compound. In some cases, hydorgen can take the place of a metal. When the element that is doing the replacing is a nonmetal, it must replace another nonmetal in a compound, and the general equation becomes:
\[\ce{Y} + \ce{XZ} \rightarrow \ce{XY} + \ce{Z}\]
where \(\ce{Y}\) is a nonmetal and replaces the nonmetal \(\ce{Z}\) in the compound with \(\ce{X}\).
Metal Replacement
Magnesium is a more reactive metal than copper. When a strip of magnesium metal is placed in an aqueous solution of copper (II) nitrate, it replaces the copper. The products of the reaction are aqueous magnesium nitrate and solid copper metal.
\[\ce{Mg} \left( s \right) + \ce{Cu(NO_3)_2} \left( aq \right) \rightarrow \ce{Mg(NO_3)_2} \left( aq \right) + \ce{Cu} \left( s \right)\]
This subcategory of single-replacement reactions is called a metal replacement reaction because it is a metal that is being replaced (copper).
Many metals react easily with acids and when they do so, one of the products of the reaction is hydrogen gas. Zinc reacts with hydrochloric acid to produce aqueous zinc chloride and hydrogen (figure below).
\[\ce{Zn} \left( s \right) + 2 \ce{HCl} \left( aq \right) \rightarrow \ce{ZnCl_2} \left( aq \right) + \ce{H_2} \left( g \right)\]
Some metals are so reactive that they are capable of replacing the hydrogen in water. The products of such a reaction are the metal hydroxide and hydrogen gas. All Group 1 metals undergo this type of reaction. Sodium reacts vigorously with water to produce aqueous sodium hydroxide and hydrogen (see figure below).
\[2 \ce{Na} \left( s \right) + 2 \ce{H_2O} \left( l \right) \rightarrow 2 \ce{NaOH} \left( aq \right) + \ce{H_2} \left( g \right)\]
Nonmetal Replacement
The element chlorine reacts with an aqueous solution of sodium bromide to produce aqueous sodium chloride and elemental bromine:
\[\ce{Cl_2} \left( g \right) + 2 \ce{NaBr} \left( aq \right) \rightarrow 2 \ce{NaCl} \left( aq \right) + \ce{Br_2} \left( l \right)\]
The reactivity of the halogen group (group 17) decreases from top to bottom within the group. Fluorine is the most reactive halogen, while iodine is the least. Since chlorine is above bromine, it is more reactive than bromine and can replace it in a halogen replacement reaction.
Double Replacement Reactions
A double-replacement reaction is a reaction in which the positive and negative ions of two ionic compounds exchange places to form two new compounds. The general form of a double-replacement (also called double-displacement) reaction is:
\[\ce{AB} + \ce{CD} \rightarrow \ce{AD} + \ce{BC}\]
In this reaction, \(\ce{A}\) and \(\ce{C}\) are positively-charged cations, while \(\ce{B}\) and \(\ce{D}\) are negatively-charged anions. Double-replacement reactions generally occur between substances in aqueous solution. In order for a reaction to occur, one of the products is usually a solid precipitate, a gas, or a molecular compound such as water.
The following are all examples of double replacement reactions:
\[2 \ce{KI} \left( aq \right) + \ce{Pb(NO_3)_2} \left( aq \right) \rightarrow 2 \ce{KNO_3} \left( aq \right) + \ce{PbI_2} \left( s \right) \label{eq10}\]
\[\ce{Na_2S} \left( aq \right) + 2 \ce{HCl} \left( aq \right) \rightarrow 2 \ce{NaCl} \left( aq \right) + \ce{H_2S} \left( g \right)\]
\[\ce{HCl} \left( aq \right) + \ce{NaOH} \left( aq \right) \rightarrow \ce{NaCl} \left( aq \right) + \ce{H_2O} \left( l \right)\]