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7.10: Classifying Chemical Reactions

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     Learning Objectives
    • Classify a chemical reaction as a synthesis, decomposition, single replacement, double replacement, or a combustion reaction.
    • Predict the products of simple reactions.

    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.

    Most chemical reactions can be classified into one or more of five basic types: acid–base reactions, exchange reactions, condensation reactions (and the reverse, cleavage reactions), and oxidation–reduction reactions. The general forms of these five kinds of reactions are summarized in Table \(\PageIndex{1}\), along with examples of each. It is important to note, however, that many reactions can be assigned to more than one classification, as you will see in our discussion.

    Table \(\PageIndex{1}\): Basic Types of Chemical Reactions
    Name of Reaction General Form Examples
    Oxidation–Reduction (redox) oxidant + reductant → reduced oxidant + oxidized reductant C7H16(l) + 11O2(g) → 7CO2(g) + 8H2O(g)
    Acid–Base acid + base → salt NaOH(aq) + HNO3(aq) → NaNO3(aq) +H2O(l)
    Exchange: Single Replacement AB + C → AC + B ZnCl2(aq)+ Mg(s) → MgCl2(aq)+ Zn(s)
    Exchange: Double Replacement AB + CDAD + CB BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq)
    Combination (Synthesis) 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. Oxidation–reduction reactions, in which there is a net transfer of electrons from one atom to another, and condensation reactions are discussed in this section. Acid–base reactions are one kind of exchange reaction—the formation of an insoluble salt, such as barium sulfate, when solutions of two soluble salts are mixed together.

    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:

    \[\ce{A} + \ce{B} \rightarrow \ce{AB} \nonumber \]

    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:

    \[2 \ce{Fe} \left( s \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{FeO} \left( s \right) \nonumber \]

    \[4 \ce{Fe} \left( s \right) + 3 \ce{O_2} \left( g \right) \rightarrow 2 \ce{Fe_2O_3} \left( s \right) \nonumber \]

    Example \(\PageIndex{1}\): Combustion of Solid Potassium

    Potassium is a very reactive alkali metal that must be stored under oil in order to prevent it from reacting with air. Write the balanced chemical equation for the combination reaction of potassium with oxygen.

    Solutions to Example 7.10.1
    Steps Example Solution
    Plan the problem. Make sure formulas of all reactants and products are correct before balancing the equation. Oxygen gas is a diatomic molecule. Potassium oxide is an ionic compound and so its formula is constructed by the crisscross method. Potassium as an ion becomes \(\ce{K^+}\), while the oxide ion is \(\ce{O^{2-}}\).

    The skeleton (unbalanced) equation:

    \[\ce{K} \left( s \right) + \ce{O_2} \left( g \right) \rightarrow \ce{K_2O} \left( s \right) \nonumber \]

    The equation is then easily balanced with coefficients.

    \[4 \ce{K} \left( s \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{K_2O} \left( s \right) \nonumber \]

    Think about your result. Formulas are correct and the resulting combination reaction is balanced.

    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:

    \[2 \ce{CO} \left( g \right) + \ce{O_2} \left( g \right) \rightarrow 2 \ce{CO_2} \left( g \right) \nonumber \]

    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:

    \[\ce{CaO} \left( s \right) + \ce{H_2O} \left( l \right) \rightarrow \ce{Ca(OH)_2} \left( aq \right) \nonumber \]

    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.

    \[\ce{SO_3} \left( g \right) + \ce{H_2O} \left( l \right) \rightarrow \ce{H_2SO_4} \left( aq \right) \nonumber \]

    Figure \(\PageIndex{1}\): Acid rain has severe consequences on both natural and manmade objects. Acid rain degrades marble statues like the one on the left (A). The trees in the forest on the right (B) have been killed by acid rain.
    Exercise \(\PageIndex{1}\)
    1. Write the chemical equation for the synthesis of silver bromide, \(\ce{AgBr}\).
    2. Predict the products for the following reaction: \(\ce{CO_2} \left( g \right) + \ce{H_2O} \left( l \right)\)
    Answer a:
    \(2 \ce{Ag} + \ce{Br_2} \rightarrow 2 \ce{AgBr}\)
    Answer b:
    \(\ce{CO_2} \left( g \right) + \ce{H_2O} \left( l \right) \rightarrow \ce{H_2CO_3}\) ​​​​​​

    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:

    \[\ce{AB} \rightarrow \ce{A} + \ce{B} \nonumber \]

    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:

    \[2 \ce{HgO} \left( s \right) \rightarrow 2 \ce{Hg} \left( l \right) + \ce{O_2} \left( g \right) \nonumber \]

    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:

    \[\ce{CaCO_3} \left( s \right) \rightarrow \ce{CaO} \left( s \right) + \ce{CO_2} \left( g \right) \nonumber \]

    Metal hydroxides decompose on heating to yield metal oxides and water. Sodium hydroxide decomposes to produce sodium oxide and water:

    \[2 \ce{NaOH} \left( s \right) \rightarrow \ce{Na_2O} \left( s \right) + \ce{H_2O} \left( g \right) \nonumber \]

    Some unstable acids decompose to produce nonmetal oxides and water. Carbonic acid decomposes easily at room temperature into carbon dioxide and water:

    \[\ce{H_2CO_3} \left( aq \right) \rightarrow \ce{CO_2} \left( g \right) + \ce{H_2O} \left( l \right) \nonumber \]

    Example \(\PageIndex{2}\): Electrolysis of Water

    When an electric current is passed through pure water, it decomposes into its elements. Write a balanced equation for the decomposition of water.

    Solutions to Example 7.10.2
    Steps Example Solution
    Plan the problem. Water is a binary compound composed of hydrogen and oxygen. The hydrogen and oxygen gases produced in the reaction are both diatomic molecules.

    The skeleton (unbalanced) equation:

    \[\ce{H_2O} \left( l \right) \overset{\text{elec}}{\rightarrow} \ce{H_2} \left( g \right) + \ce{O_2} \left( g \right) \nonumber \]

    Note the abbreviation "\(\text{elec}\)" above the arrow to indicate the passage of an electric current to initiate the reaction. Balance the equation.

    \[2 \ce{H_2O} \left( l \right) \overset{\text{elec}}{\rightarrow} 2 \ce{H_2} \left( g \right) + \ce{O_2} \left( g \right) \nonumber \]

    Think about your result. The products are elements and the equation is balanced.
    Exercise \(\PageIndex{2}\)

    Write the chemical equation for the decomposition of:

    1. \(\ce{Al_2O_3}\)
    2. \(\ce{Ag_2S}\)
    Answer a
    \(2 \ce{Al_2O_3} \rightarrow 4 \ce{Al} + 3 \ce{O_2}\)
    Answer b
    \(\ce{Ag_2S} \rightarrow 2 \ce{Ag} + \ce{S}\)

    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:

    \[\ce{A} + \ce{BC} \rightarrow \ce{AC} + \ce{B} \nonumber \]

    In this general reaction, element \(\ce{A}\) is a metal and replaces element \(\ce{B}\), also a metal, in the compound. When the element that is doing the replacing is a nonmetal, it must replace another nonmetal in a compound, and the general equation becomes:

    \[\ce{Y} + \ce{XZ} \rightarrow \ce{XY} + \ce{Z} \nonumber \]

    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.

    \[\ce{Mg} \left( s \right) + \ce{Cu(NO_3)_2} \left( aq \right) \rightarrow \ce{Mg(NO_3)_2} \left( aq \right) + \ce{Cu} \left( s \right) \nonumber \]

    This subcategory of single-replacement reactions is called a metal replacement reaction because it is a metal that is being replaced (copper).

    Hydrogen Replacement

    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).

    \[\ce{Zn} \left( s \right) + 2 \ce{HCl} \left( aq \right) \rightarrow \ce{ZnCl_2} \left( aq \right) + \ce{H_2} \left( g \right) \nonumber \]

    In a hydrogen replacement reaction, the hydrogen in the acid is replaced by an active metal. 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).

    \[2 \ce{Na} \left( s \right) + 2 \ce{H_2O} \left( l \right) \rightarrow 2 \ce{NaOH} \left( aq \right) + \ce{H_2} \left( g \right) \nonumber \]

    Sodium reacts with water to produce hydrogen.
    Figure \(\PageIndex{2}\): (First image) Zinc metal reacts with hydrochloric acid to give off hydrogen gas in a single-displacement reaction. (Second image) Sodium metal reacts vigorously with water, giving off hydrogen gas. A large piece of sodium will often generate so much heat that the hydrogen will ignite.

    Halogen Replacement

    The element chlorine reacts with an aqueous solution of sodium bromide to produce aqueous sodium chloride and elemental bromine:

    \[\ce{Cl_2} \left( g \right) + 2 \ce{NaBr} \left( aq \right) \rightarrow 2 \ce{NaCl} \left( aq \right) + \ce{Br_2} \left( l \right) \nonumber \]

    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.

    Example \(\PageIndex{3}\)

    What are the products of the reaction between solid aluminum (\(\ce{Al}\)) and iron (III) oxide (\(\ce{Fe_2O_3}\))?

    Solutions to Example 7.10.3
    Steps Example Solution
    Plan the problem. To predict the products, we need to know that aluminum will replace iron and form aluminum oxide (the metal will replace the metal ion in the compound). Aluminum has a charge of \(+3\) and oxygen has a charge of \(-2\). The compound formed between aluminum and oxygen, therefore, will be \(\ce{Al_2O_3}\). Since iron is replaced in the compound by aluminum, the iron will now be the single element in the products.

    The unbalanced equation will be:

    \[\ce{Al} + \ce{Fe_2O_3} \rightarrow \ce{Al_2O_3} + \ce{Fe} \nonumber \]

    and the balanced equation will be:

    \[2 \ce{Al} + \ce{Fe_2O_3} \rightarrow \ce{Al_2O_3} + 2 \ce{Fe} \nonumber \]

    Think about your result. This is a single replacement reaction, and when balanced the coefficients accurately reflect that the iron and aluminum have the same charge in this reaction.
    Exercise \(\PageIndex{3}\)
    1. Write the chemical equation for the single replacement reaction between zinc solid and lead (II) nitrate solution to produce zinc nitrate solution and solid lead. (Note that zinc forms ions with a \(+2\) charge.)
    2. Predict the products for the following reaction: \(\ce{Fe} + \ce{CuSO_4}\). (In this reaction, assume iron forms ions with a \(+2\) charge.)
    Answer a
    \(\ce{Zn} + \ce{Pb(NO_3)_2} \rightarrow \ce{Pb} + \ce{Zn(NO_3)_2}\)
    Answer b
    \(\ce{Fe} + \ce{CuSO_4} \rightarrow \ce{Cu} + \ce{FeSO_4}\)

    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:

    \[\ce{AB} + \ce{CD} \rightarrow \ce{AD} + \ce{BC} \nonumber \]

    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.

    Formation of a Precipitate

    A precipitate forms in a double-replacement reaction when the cations from one of the reactants combine with the anions from the other reactant to form an insoluble ionic compound. When aqueous solutions of potassium iodide and lead (II) nitrate are mixed, the following reaction occurs:

    \[2 \ce{KI} \left( aq \right) + \ce{Pb(NO_3)_2} \left( aq \right) \rightarrow 2 \ce{KNO_3} \left( aq \right) + \ce{PbI_2} \left( s \right) \label{eq10} \]

    There are very strong attractive forces that occur between \(\ce{Pb^{2+}}\) and \(\ce{I^-}\) ions and the result is a brilliant yellow precipitate (Figure \(\PageIndex{3}\)). The other product of the reaction, potassium nitrate, remains soluble.

    Yellow lead (II) iodide precipitate in test tube
    Figure \(\PageIndex{3}\): Lead (II) iodide precipitates when potassium iodide is mixed with lead (II) nitrate (Equation \ref{eq10}). (CC BY-SA 3.0; PRHaney).

    Formation of a Gas

    Some double-replacement reactions produce a gaseous product which then bubbles out of the solution and escapes into the air. When solutions of sodium sulfide and hydrochloric acid are mixed, the products of the reaction are aqueous sodium chloride and hydrogen sulfide gas:

    \[\ce{Na_2S} \left( aq \right) + 2 \ce{HCl} \left( aq \right) \rightarrow 2 \ce{NaCl} \left( aq \right) + \ce{H_2S} \left( g \right) \nonumber \]

    Formation of a Molecular Compound

    Another kind of double-replacement reaction is one that produces a molecular compound as one of its products. Many examples in this category are reactions that produce water. When aqueous hydrochloric acid is reacted with aqueous sodium hydroxide, the products are aqueous sodium chloride and water:

    \[\ce{HCl} \left( aq \right) + \ce{NaOH} \left( aq \right) \rightarrow \ce{NaCl} \left( aq \right) + \ce{H_2O} \left( l \right) \nonumber \]

    Example \(\PageIndex{4}\)

    Write a complete and balanced chemical equation for the double-replacement reaction \(\ce{NaCN} \left( aq \right) + \ce{HBr} \left( aq \right) \rightarrow\) (hydrogen cyanide gas is formed).

    Solutions to Example 7.10.4
    Steps Example Solution
    Plan the problem. The production of a gas drives the reaction.

    The cations of both reactants are \(+1\) charged ions, while the anions are \(-1\) charged ions. After exchanging partners, the balanced equation is:

    \[\ce{NaCN} \left( aq \right) + \ce{HBr} \left( aq \right) \rightarrow \ce{NaBr} \left( aq \right) + \ce{HCN} \left( g \right) \nonumber \]

    Think about your result. This is a double replacement reaction. All formulas are correct and the equation is balanced.
    Exercise \(\PageIndex{4}\)

    Write a complete and balanced chemical equation for the double-replacement reaction \(\ce{(NH_4)_2SO_4} \left( aq \right) + \ce{Ba(NO_3)_2} \left( aq \right) \rightarrow\) (a precipitate of barium sulfate forms).

    Answer a:
    \[\ce{(NH_4)_2SO_4} \left( aq \right) + \ce{Ba(NO_3)_2} \left( aq \right) \rightarrow 2 \ce{NH_4NO_3} \left( aq \right) + \ce{BaSO_4} \left( s \right) \nonumber \]

    Occasionally, a reaction will produce both a gas and a molecular compound. The reaction of a sodium carbonate solution with hydrochloric acid produces aqueous sodium chloride, carbon dioxide gas, and water:

    \[\ce{Na_2CO_3} \left( aq \right) + 2 \ce{HCl} \left( aq \right) \rightarrow 2 \ce{NaCl} \left( aq \right) + \ce{CO_2} \left( g \right) + \ce{H_2O} \left( l \right) \nonumber \]

    7.10: Classifying Chemical Reactions is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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