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11.2: Types of Chemical Reactions

Created by: CK-12

Lesson Objectives

  • Be able to classify a chemical reaction as a combination, decomposition, single replacement, double replacement, or combustion reaction.
  • Be able to predict the products when given a set of reactants for a given chemical process.
  • Explain the concept of solubility and the process of precipitation.
  • Use solubility information to predict whether or not a given substance is soluble in water.
  • Use the general solubility rules to predict chemical behavior.
  • Be able to write molecular, ionic, and net ionic equations for a given chemical process.

Lesson Vocabulary

  • combination reaction: A reaction where two or more chemical species combine to produce a single new compound.
  • decomposition reaction: A reaction where a single chemical species breaks down to produce two or more new chemical species.
  • single replacement reaction: Occurs when one chemical species (often a single element) replaces a portion of another compound to produce two new products.
  • double replacement reaction: Occurs when the cations from the original two ionic compounds trade anions to make two new ionic compounds.
  • molecular equation: An equation that shows all ionic components as neutral compounds, but the ones that are dissolved in water are denoted with "(aq)."
  • ionic equation: A chemical equation in which the various reaction components are represented as they actually exist in the reaction, for example, as individual ions.
  • spectator ion: Ions that are present in solution but do not participate in the overall reaction.
  • net ionic equation: The simplified ionic equation in which all of the spectator ions are cancelled out.
  • combustion: Occurs when a hydrocarbon reacts in the presence of oxygen to produce water and carbon dioxide.

Check Your Understanding

Study the Figure below, which depicts the mass change that occurs when steel wool burns in air.

Mass changes for steel wool burning in air

  1. What happens to the mass of the steel wool as the reaction proceeds?
  2. Given that mass must be conserved in chemical reactions (it cannot come from nowhere), what might be your explanation for the change in the mass of the steel wool?
  3. How might mass changes such as this help us identify and categorize a given chemical process?

Introduction

The video above at http://www.youtube.com/watch?v=_Y1alDuXm6A (1:12) shows the decomposition of mercury(II) oxide into liquid mercury and oxygen gas. This reaction was an important one in the history of chemistry, because it helped early chemists to understand the relationship between reactants and products.

In the last lesson, we began investigating how a chemical equation can represent a given chemical reaction. In this lesson, we are going to study the ways in which chemical reactions are classified. There are literally thousands of chemical reactions that take place every day in our lives. Some reactions take place in the atmosphere, such as the combustion of fossil fuels. Others occur in solution, like the reactions responsible for photosynthesis or the reactions that break down our food to give us energy. Chemical reactions can take place in a variety of environments. Reactions happen on the sea floor, in our cells, and in the upper atmosphere. As we look at chemical reactions, we notice some commonalities and trends. When we studied the elements, we saw characteristics that allowed us to categorize them by family. There are also various ways to categorize chemical reactions. Some reactions produce heat, while others consume it. Some reactions are spontaneous, while others are not. Some reactions happen in nanoseconds, while others happen over longer spans of time. Some produce electricity, some emit light, and some release gaseous products. The products of chemical reactions tell us a lot about the chemistry of the process. In the above video, we see mercury(II) oxide decomposing into elemental mercury and oxygen gas. Decomposition was one of the first reaction types to be identified by chemists. Decomposition is one type of reaction you'll learn about in this lesson.

Combination Reactions

The first type of reaction that we will investigate is the combination reaction, which is sometimes also referred to as a synthesis reaction. In combination reactions, two or more chemical species combine to produce a single new compound. A generic combination reaction might have the following form:

A+B \rightarrow C

Substances in all states of matter can participate in combination reactions. For example, oxygen in the air can react with iron to produce rust. Rusting is a common occurrence, especially in regions of the world where precipitation is relatively high. Although rust tends to be a mixture of compounds, its primary component is iron(III) oxide (Fe2O3). Rusting is generally a very slow process, but when the iron has a very high surface area, as in the case of steel wool, it can happen at a much faster rate, as shown in the following video:

http://www.youtube.com/watch?v=5MDH92VxPEQ

The balanced chemical equation for this process is shown below:

4\text{Fe}(s)+3\text{O}_{2}(g) \rightarrow 2\text{Fe}_2\text{O}_{3}(s)

Decomposition Reactions

A decomposition reaction is the exact opposite of a combination reaction. In decomposition reactions, a single chemical species breaks down to produce two or more new chemical species. A generic decomposition reaction might take the following form:

C \rightarrow A+B

Again, substances in all states of matter commonly participate in decomposition reactions. For example, hydrogen peroxide will decompose over time to produce water and oxygen gas according to the following equation:

2\text{H}_2\text{O}_{2}(l) \rightarrow 2\text{H}_2\text{O}(l)+\text{O}_{2}(g)

Another common type of decomposition reaction involves the process of electrolysis, in which an electrical current is passed through a substance to break apart a compound. One example of a decomposition reaction requiring the use of electrolysis is the decomposition of molten sodium chloride, as shown by the following equation:

2\text{NaCl}(s) \rightarrow 2\text{Na}(s)+\text{Cl}_{2}(g)

Single Replacement Reactions

A single replacement reaction (sometimes called a single displacement reaction) occurs when one chemical species (often a single element) replaces a portion of another compound to produce two new products. The general form of a single replacement reaction is shown below:

AB+C \rightarrow AC+B

Two common types of single replacement reactions involve pure metals reaction with aqueous solutions of either an acid or an ionic compound. When a reactive metal is placed in an acid solution, the following reaction is likely to occur:

Metal + acid → ionic solution + hydrogen gas

An example of this would be the reaction between zinc and hydrochloric acid, which produces zinc chloride and hydrogen gas. Here is an image of this reaction:

Zinc metal reacting with a solution of hydrochloric acid

The balanced chemical equation for this single replacement reaction is shown below:

\text{Zn}(s)+2\text{HCl}(aq) \rightarrow \text{ZnCl}_{2}(aq)+\text{H}_{2}(g)

Another type of single replacement reaction involves a solid metal replacing the metal cation in an ionic compound that has been dissolved in water. If the solid metal is more reactive than the dissolved metal cations, the following type of reaction can occur:

Metal + ionic solution → different metal + different ionic solution

A common example of this reaction is when iron is replaced by the more reactive zinc metal. The balanced chemical equation for this process is shown below.

\text{Zn}(s)+\text{FeSO}_{4}(aq) \rightarrow \text{Fe}(s)+\text{ZnSO}_{4}(aq)

Double Replacement Reactions

Double replacement reactions typically include two water-soluble salts that react with one another in solution. The general form of a double replacement reaction would look something like the following:

AB+CD \rightarrow AD+CB

In double replacement reactions, the cations from the original two ionic compounds trade anions to make two new ionic compounds. In general, at least one of the new compounds must precipitate (form an insoluble solid) for us to conclude that a reaction has occurred. An example of such a process is shown below with the double replacement reaction between solutions of potassium iodide and lead(II) nitrate.

A double replacement reaction is used to form lead(II) iodide. The reactants shown here are colorless solutions of potassium iodide and potassium nitrate. When combined, these produce a yellow precipitate of lead(II) iodide.

At the molecular level, our model for the way in which a precipitate forms can be described in an animation:

http://www.crescent.edu.sg/crezlab/webpages/PptReaction_PbI2.htm

Representing Ionic Reactions as Chemical Equations

For reactions that involve ions dissolved in water, there are several different ways to express the overall process as a chemical equation. For example, the overall molecular equation shows all ionic components as neutral compounds, but the ones that are dissolved in water are denoted with "(aq)." Note that the ionic substances do not exist as molecules, but we write them out as though they were. In the following example, two water-soluble compounds trade partners to produce one dissolved ionic compound and one solid precipitate:

AB(aq)+CD(aq) \rightarrow AD(aq)+CB(s)

In reality, the aqueous substances do not exist as molecules or ionic crystal lattices. Instead, the individual ions are dissolved and distributed throughout the solution. If the reaction above were written as an ionic equation, it would look something like the following:

A^+(aq)+B^-(aq)+C^+(aq)+D^-(aq) \rightarrow A^+(aq)+D^-(aq)+CB(s)

In this example, the various reaction components are presented in a form that is closer to the way they actually exist during the reaction. The aqueous components are separated into ions, and the precipitate is found as a combined solid. We are assuming in this example that A and C form cations with a charge of 1+, while B and D form anions with a charge of 1-. In real examples, we would look at which group each element is found in on the periodic table to determine its likely charge.

Notice that in the ionic equation, A+ and D- were unchanged over the course of the reaction; they exist as aqueous ions on both the reactant and product sides. In other words, these species did not experience any net change. Ions that are present in solution but do not participate in the overall reaction are known as spectator ions. The ionic equation can be simplified to the net ionic equation by canceling out all the spectator ions.

\cancel{A^+(aq)}+B^-(aq)+C^+(aq)+\cancel{D^-(aq)} & \rightarrow \cancel{A^+(aq)}+\cancel{D^-(aq)}+CB(s) \\B^-(aq)+C^+(aq) & \rightarrow CB(s)

Let's look at these three types of equations again using a real example. If we were to mix aqueous solutions of potassium iodide and lead(II) nitrate, lead(II) iodide would precipitate as a solid, and potassium nitrate would remain dissolved. This can be represented by any of the three following equations:

Molecular Equation

2\text{KI}(aq)+\text{Pb}(\text{NO}_3)_{2}(aq) \rightarrow 2\text{KNO}_{3}(aq)+\text{PbI}_{2}(s)

Ionic Equation

2\text{K}^+(aq)+2\text{I}^-(aq)+\text{Pb}^{2+}(aq)+2\text{NO}^-_{3}(aq) \rightarrow 2\text{K}^+(aq)+2\text{NO}^-_{3}(aq)+\text{PbI}_{2}(s)

Net Ionic Equation

\text{Pb}^{2+}(aq)+2\text{I}^-(aq) \rightarrow \text{PbI}_{2}(s)

Predicting Solubility of Ionic Compounds

How do we determine which ions are likely to form an insoluble precipitate and which will remain dissolved in water? By combining various ionic solutions, chemists have come up with some general guidelines for whether a given cation-anion pairing is likely to be soluble or insoluble in water. It should be noted that such an approach is an oversimplification. Each compound has its own solubility value, so two "soluble" compounds might have very different abilities to dissolve in water. Additionally, even "insoluble" salts can dissolve in water to a very limited extent. We will take a more quantitative approach to solubility in the chapter on solutions. However, qualitative rules like the ones in the Table below are useful for predicting whether a precipitate is likely to form when combining moderate amounts of specific cations and anions.

Solubility Properties to Predict Products of Chemical Reactions
Type of Particle Soluble Insoluble
Common Cations Alkali metal cation (Li+, Na+, K+, Rb+, or Cs+) or the NH4+ cation
Common Anions ClO4- and NO3- compounds
Halides Most Cl-, Br-, and I- compounds Compounds that include the Ag+, Pb2+, or Hg22+ cations
Sulfates Most SO42- compounds PbSO4, Ag2SO4, Hg2SO4, CaSO4, SrSO4, and BaSO4
Sulfides Compounds with NH4+ or a metal from group IA or IIA as cation Most S2- compounds
Hydroxides Compounds with NH4+, Ba2+, or a metal from group IA as cation Most OH- compounds
Carbonates, phosphates, and sulfites Compounds with NH4+ or a metal from group IA as cation Most CO32- and PO43-, and SO32- compounds

Combustion

Combustion occurs when a hydrocarbon reacts in the presence of oxygen to produce water and carbon dioxide. These reactions are very exothermic, which means that they produce a large amount of heat. Combustion reactions are quite common in our everyday lives, such as the burning of gasoline to fuel a car. The chemical equation for a combustion reaction has the following generic form:

\text{C}_{x}\text{H}_{y}+\text{O}_2 \rightarrow \text{H}_2\text{O}+\text{CO}_2

Combustion reaction of a marshmallow (sucrose) and wood (cellulose).

The process of cellular respiration can be thought of as a highly controlled version of a combustion reaction. We do not literally burn hydrocarbons in our body, but the overall reactants and products are the same. Hydrocarbons, such as sucrose (C12H22O11), are combined with oxygen in a series of enzymatic steps to product water, carbon dioxide, and energy, which is stored in the form of reactive molecules. The unbalanced chemical equation for this overall process is shown below:

\text{C}_{12}\text{H}_{22}\text{O}_{11}+\text{O}_2 \rightarrow \text{CO}_2+\text{H}_2\text{O}

Lesson Summary

  • Combination reactions occur when two or more reactants combine to produce a single compound.
  • Decomposition reactions involve one compound decomposing into two or more products.
  • Single replacement reactions occur when one reactant replaces part of another compound to form new substances.
  • A common type of double replacement reaction occurs when two ionic reactants exchange anions, making two new ionic compounds. The precipitation of a solid is a common result for this type of reaction.
  • Combustion reactions involve the reaction of a hydrocarbon with oxygen gas to produce water and carbon dioxide.

Lesson Review Questions

  1. Categorize the following chemical reactions as single replacement, double replacement, combustion, combination, or decomposition.
    1. Equimolar (having the same number of moles) solutions of silver nitrate and potassium chloride are mixed to produce solid silver chloride and aqueous potassium nitrate.
    2. Magnesium metal is added to hydrochloric acid to produce hydrogen gas and aqueous magnesium chloride.
    3. Ethanol is burned in air to produce water and carbon dioxide gas.
    4. Water is electrolyzed to produce hydrogen and oxygen gas.
    5. Hydrogen gas and oxygen gas are ignited to produce water.
  2. Write the balanced chemical equation for the following combination and decomposition reactions.
    1. Magnesium carbonate is heated strongly to produce magnesium oxide and carbon dioxide gas.
    2. Hydrogen peroxide decomposes to produce water and oxygen gas.
    3. Solid potassium chlorate is heated in the presence of manganese dioxide as a catalyst to produce potassium chloride and oxygen gas. (Catalysts speed up reactions but are not expressed in the overall balanced equation)
    4. Molten aluminum oxide is electrolyzed using inert (non-reactive) electrodes to produce aluminum metal and oxygen gas.
  3. Write the balanced chemical equations for the following replacement reactions:
    1. Zinc metal is added to a solution of iron(II) sulfate.
    2. Equimolar solutions of lead(II) nitrate and sodium chloride are mixed to produce solid lead(II) chloride and aqueous sodium nitrate.
    3. Solutions of potassium phosphate and zinc nitrate are mixed.
  4. Write the balanced chemical equations for the following combustion reactions.
    1. Propane (C3H8) is ignited in air to produce water and carbon dioxide gas.
    2. Methanol(CH4O) is ignited in air to produce water and carbon dioxide gas.
    3. Ethanol (C2H5OH) is burned in air.
  5. Write the molecular equation, ionic equation, and net ionic equation for each of the following double replacement reactions.
    1. Silver nitrate reacts with potassium iodide to produce potassium nitrate and silver iodide.
    2. Silver nitrate reacts with iron(III) chloride to produce iron(III) nitrate and silver chloride.
    3. Lead(II) nitrate reacts with potassium iodide to produce potassium nitrate and lead(II) iodide.
    4. Iron(III) chloride reacts with lead(II) nitrate to produce lead(II) chloride and iron(III) nitrate.
    5. Calcium chloride reacts with sodium hydroxide to produce calcium hydroxide and sodium chloride.
  6. Would it be possible to have a double precipitate formed for a double replacement process? Can you write an equation where a double precipitate forms?
  7. What is meant when we describe a compound as (aq) or (s)? Explain the similarities and differences between these terms.
  8. Write the balanced chemical equation for the combination reaction in which hydrogen and oxygen gases react explosively to produce water. (Remember that hydrogen and oxygen exist as diatomic gases in their most common elemental form.)
  9. Write the balanced chemical equation for the reaction that occurs when a piece of aluminum metal is placed in a solution of silver nitrate.
  10. Using the solubility rules given above, predict whether or not the following compounds are soluble or insoluble in water.
    1. Potassium nitrate
    2. Lead(II) chloride
    3. Barium sulfate
    4. Aluminum sulfide
    5. Calcium carbonate

Further Reading / Supplemental Links

Points to Consider

  1. In an earlier section, we discussed the origins of the chemical recipe for gunpowder, one of the earliest chemical formulas to be described. The recipe for gun powder is 75 percent potassium nitrate, 15 percent charcoal, and 10 percent sulfur. How might one measure out these amounts in a predictable and reliable way?
  2. So far, we have discussed the characteristics of a variety of reactions. However, we have spent little time discussing how we might measure and calculate amounts of reactants and products. The steel wool reaction is as follows: 4\text{Fe}_{(s)}+3\text{O}_{2(g)} \rightarrow 2\text{Fe}_2\text{O}_{3(s)}. How might you measure the amounts of each reactant used and the product that forms?
  3. In the chemical reactions that we have already studied, we have assumed that all reactants are transformed into products (the reaction "goes to completion"). Are there reactions that do not go to completion? How do you know whether you will have reactants left over?
  4. What are some factors that control whether or not a chemical reaction takes place?

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Sep 09, 2013

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Nov 24, 2014
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