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Introductory Chemistry - 1st Canadian Edition: Types of Chemical Reactions: Single- and Double-Displacement Reactions

Introductory Chemistry - 1st Canadian Edition
Types of Chemical Reactions: Single- and Double-Displacement Reactions
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table of contents
  1. Cover
  2. Title Page
  3. Copyright
  4. Table Of Contents
  5. Acknowledgments
  6. Dedication
  7. About BCcampus Open Education
  8. Chapter 1. What is Chemistry
    1. Some Basic Definitions
    2. Chemistry as a Science
  9. Chapter 2. Measurements
    1. Expressing Numbers
    2. Significant Figures
    3. Converting Units
    4. Other Units: Temperature and Density
    5. Expressing Units
    6. End-of-Chapter Material
  10. Chapter 3. Atoms, Molecules, and Ions
    1. Acids
    2. Ions and Ionic Compounds
    3. Masses of Atoms and Molecules
    4. Molecules and Chemical Nomenclature
    5. Atomic Theory
    6. End-of-Chapter Material
  11. Chapter 4. Chemical Reactions and Equations
    1. The Chemical Equation
    2. Types of Chemical Reactions: Single- and Double-Displacement Reactions
    3. Ionic Equations: A Closer Look
    4. Composition, Decomposition, and Combustion Reactions
    5. Oxidation-Reduction Reactions
    6. Neutralization Reactions
    7. End-of-Chapter Material
  12. Chapter 5. Stoichiometry and the Mole
    1. Stoichiometry
    2. The Mole
    3. Mole-Mass and Mass-Mass Calculations
    4. Limiting Reagents
    5. The Mole in Chemical Reactions
    6. Yields
    7. End-of-Chapter Material
  13. Chapter 6. Gases
    1. Pressure
    2. Gas Laws
    3. Other Gas Laws
    4. The Ideal Gas Law and Some Applications
    5. Gas Mixtures
    6. Kinetic Molecular Theory of Gases
    7. Molecular Effusion and Diffusion
    8. Real Gases
    9. End-of-Chapter Material
  14. Chapter 7. Energy and Chemistry
    1. Formation Reactions
    2. Energy
    3. Stoichiometry Calculations Using Enthalpy
    4. Enthalpy and Chemical Reactions
    5. Work and Heat
    6. Hess’s Law
    7. End-of-Chapter Material
  15. Chapter 8. Electronic Structure
    1. Light
    2. Quantum Numbers for Electrons
    3. Organization of Electrons in Atoms
    4. Electronic Structure and the Periodic Table
    5. Periodic Trends
    6. End-of-Chapter Material
  16. Chapter 9. Chemical Bonds
    1. Lewis Electron Dot Diagrams
    2. Electron Transfer: Ionic Bonds
    3. Covalent Bonds
    4. Other Aspects of Covalent Bonds
    5. Violations of the Octet Rule
    6. Molecular Shapes and Polarity
    7. Valence Bond Theory and Hybrid Orbitals
    8. Molecular Orbitals
    9. End-of-Chapter Material
  17. Chapter 10. Solids and Liquids
    1. Properties of Liquids
    2. Solids
    3. Phase Transitions: Melting, Boiling, and Subliming
    4. Intermolecular Forces
    5. End-of-Chapter Material
  18. Chapter 11. Solutions
    1. Colligative Properties of Solutions
    2. Concentrations as Conversion Factors
    3. Quantitative Units of Concentration
    4. Colligative Properties of Ionic Solutes
    5. Some Definitions
    6. Dilutions and Concentrations
    7. End-of-Chapter Material
  19. Chapter 12. Acids and Bases
    1. Acid-Base Titrations
    2. Strong and Weak Acids and Bases and Their Salts
    3. Brønsted-Lowry Acids and Bases
    4. Arrhenius Acids and Bases
    5. Autoionization of Water
    6. Buffers
    7. The pH Scale
    8. End-of-Chapter Material
  20. Chapter 13. Chemical Equilibrium
    1. Chemical Equilibrium
    2. The Equilibrium Constant
    3. Shifting Equilibria: Le Chatelier’s Principle
    4. Calculating Equilibrium Constant Values
    5. Some Special Types of Equilibria
    6. End-of-Chapter Material
  21. Chapter 14. Oxidation and Reduction
    1. Oxidation-Reduction Reactions
    2. Balancing Redox Reactions
    3. Applications of Redox Reactions: Voltaic Cells
    4. Electrolysis
    5. End-of-Chapter Material
  22. Chapter 15. Nuclear Chemistry
    1. Units of Radioactivity
    2. Uses of Radioactive Isotopes
    3. Half-Life
    4. Radioactivity
    5. Nuclear Energy
    6. End-of-Chapter Material
  23. Chapter 16. Organic Chemistry
    1. Hydrocarbons
    2. Branched Hydrocarbons
    3. Alkyl Halides and Alcohols
    4. Other Oxygen-Containing Functional Groups
    5. Other Functional Groups
    6. Polymers
    7. End-of-Chapter Material
  24. Chapter 17. Kinetics
    1. Factors that Affect the Rate of Reactions
    2. Reaction Rates
    3. Rate Laws
    4. Concentration–Time Relationships: Integrated Rate Laws
    5. Activation Energy and the Arrhenius Equation
    6. Reaction Mechanisms
    7. Catalysis
    8. End-of-Chapter Material
  25. Chapter 18. Chemical Thermodynamics
    1. Spontaneous Change
    2. Entropy and the Second Law of Thermodynamics
    3. Measuring Entropy and Entropy Changes
    4. Gibbs Free Energy
    5. Spontaneity: Free Energy and Temperature
    6. Free Energy under Nonstandard Conditions
    7. End-of-Chapter Material
  26. Appendix A: Periodic Table of the Elements
  27. Appendix B: Selected Acid Dissociation Constants at 25°C
  28. Appendix C: Solubility Constants for Compounds at 25°C
  29. Appendix D: Standard Thermodynamic Quantities for Chemical Substances at 25°C
  30. Appendix E: Standard Reduction Potentials by Value
  31. Glossary
  32. About the Authors
  33. Versioning History

Types of Chemical Reactions: Single- and Double-Displacement Reactions

Learning Objectives

  1. Recognize chemical reactions as single-replacement reactions and double-replacement reactions.
  2. Use the periodic table, an activity series, or solubility rules to predict whether single-replacement reactions or double-replacement reactions will occur.

Up to now, we have presented chemical reactions as a topic, but we have not discussed how the products of a chemical reaction can be predicted. Here we will begin our study of certain types of chemical reactions that allow us to predict what the products of the reaction will be.

A single-replacement reaction is a chemical reaction in which one element is substituted for another element in a compound, generating a new element and a new compound as products. For example:

2HCl(aq) + Zn(s) → ZnCl2(aq) + H2(g)

is an example of a single-replacement reaction. The hydrogen atoms in HCl are replaced by Zn atoms, and in the process a new element—hydrogen—is formed. Another example of a single-replacement reaction is:

2NaCl(aq) + F2(g) → 2NaF(s) + Cl2(g)

Here the negatively charged ion changes from chloride to fluoride. A typical characteristic of a single-replacement reaction is that there is one element as a reactant and another element as a product.

Not all proposed single-replacement reactions will occur between two given reactants. This is most easily demonstrated with fluorine, chlorine, bromine, and iodine. Collectively, these elements are called the halogens and are in the next-to-last column on the periodic table (see Figure 4.1 “Halogens on the Periodic Table”). The elements on top of the column will replace the elements below them on the periodic table but not the other way around. Thus, the reaction represented by:

CaI2(s) + Cl2(g) → CaCl2(s) + I2(s)

This reaction will occur, but the reaction

CaF2(s) + Br2(ℓ) → CaBr2(s) + F2(g)

will not because bromine is below fluorine on the periodic table. This is just one of many ways the periodic table helps us understand chemistry.

Hydrogen, fluorine, chlorine, bromine, iodine, and astatine are halogens.
Figure 4.1 “Halogens on the Periodic Table.” The halogens are the elements in the next-to-last column on the periodic table.

Example 4.2

Problems

Will a single-replacement reaction occur? If so, identify the products.

  1. MgCl2 + I2 → ?
  2. CaBr2 + F2 → ?

Solutions

  1. Because iodine is below chlorine on the periodic table, a single-replacement reaction will not occur.
  2. Because fluorine is above bromine on the periodic table, a single-replacement reaction will occur, and the products of the reaction will be CaF2 and Br2.

Test Yourself

Will a single-replacement reaction occur? If so, identify the products.

FeI2 + Cl2 → ?

Answer

Yes; FeCl2 and I2

Chemical reactivity trends are easy to predict when replacing anions in simple ionic compounds—simply use their relative positions on the periodic table. However, when replacing the cations, the trends are not as straightforward. This is partly because there are so many elements that can form cations; an element in one column on the periodic table may replace another element nearby, or it may not. A list called the activity series does the same thing the periodic table does for halogens: it lists the elements that will replace elements below them in single-replacement reactions. A simple activity series is shown below.

Activity Series for Cation Replacement in Single-Replacement Reactions

  • Li
  • K
  • Ba
  • Sr
  • Ca
  • Na
  • Mg
  • Al
  • Mn
  • Zn
  • Cr
  • Fe
  • Ni
  • Sn
  • Pb
  • H2
  • Cu
  • Hg
  • Ag
  • Pd
  • Pt
  • Au

Using the activity series is similar to using the positions of the halogens on the periodic table. An element on top will replace an element below it in compounds undergoing a single-replacement reaction. Elements will not replace elements above them in compounds.

Example 4.3

Problems

Use the activity series to predict the products, if any, of each equation.

  1. FeCl2 + Zn → ?
  2. HNO3 + Au → ?

Solutions

  1. Because zinc is above iron in the activity series, it will replace iron in the compound. The products of this single-replacement reaction are ZnCl2 and Fe.
  2. Gold is below hydrogen in the activity series. As such, it will not replace hydrogen in a compound with the nitrate ion. No reaction is predicted.

Test Yourself

Use the activity series to predict the products, if any, of this equation.

AlPO4 + Mg → ?

Answer

Mg3(PO4)2 and Al

A double-replacement reaction occurs when parts of two ionic compounds are exchanged, making two new compounds. A characteristic of a double-replacement equation is that there are two compounds as reactants and two different compounds as products. An example is:

CuCl2(aq) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2AgCl(s)

There are two equivalent ways of considering a double-replacement equation: either the cations are swapped, or the anions are swapped. (You cannot swap both; you would end up with the same substances you started with.) Either perspective should allow you to predict the proper products, as long as you pair a cation with an anion and not a cation with a cation or an anion with an anion.

Example 4.4

Problem

Predict the products of this double-replacement equation: BaCl2 + Na2SO4 → ?

Solution

Thinking about the reaction as either switching the cations or switching the anions, we would expect the products to be BaSO4 and NaCl.

Test Yourself

Predict the products of this double-replacement equation: KBr + AgNO3 → ?

Answer

KNO3 and AgBr

Predicting whether a double-replacement reaction occurs is somewhat more difficult than predicting a single-replacement reaction. However, there is one type of double-replacement reaction that we can predict: the precipitation reaction. A precipitation reaction occurs when two ionic compounds are dissolved in water and form a new ionic compound that does not dissolve; this new compound falls out of solution as a solid precipitate. The formation of a solid precipitate is the driving force that makes the reaction proceed.

To judge whether double-replacement reactions will occur, we need to know what kinds of ionic compounds form precipitates. For this, we use solubility rules, which are general statements that predict which ionic compounds dissolve (are soluble) and which do not (are not soluble or insoluble). Tables 4.1a and b “Some Useful Solubility Rules” list some general solubility rules. We need to consider each ionic compound (both the reactants and the possible products) in light of the solubility rules in Table 4.1a and b. If a compound is soluble, we use the (aq) label with it, indicating it dissolves. If a compound is not soluble, we use the (s) label with it and assume that it will precipitate out of solution. If everything is soluble, then no reaction will be expected.

Table 4.1a Some Useful Solubility Rules
These compounds generally dissolve in water (are soluble):Exceptions:
All compounds of Li+, Na+, K+, Rb+, Cs+, and NH4+None
All compounds of NO3− and C2H3O2−None
Compounds of Cl−, Br−, I−Ag+, Hg22+, Pb2+
Compounds of SO42Hg22+, Pb2+, Sr2+, Ba2+
Table 4.1b Some More Useful Solubility Rules
These compounds generally do not dissolve in water (are insoluble):Exceptions:
Compounds of CO32− and PO43−Compounds of Li+, Na+, K+, Rb+, Cs+, and NH4+
Compounds of OH−Compounds of Li+, Na+, K+, Rb+, Cs+, NH4+, Sr2+, and Ba2+

For example, consider the possible double-replacement reaction between Na2SO4 and SrCl2. The solubility rules say that all ionic sodium compounds are soluble and all ionic chloride compounds are soluble except for Ag+, Hg22+, and Pb2+, which are not being considered here. Therefore, Na2SO4 and SrCl2 are both soluble. The possible double-replacement reaction products are NaCl and SrSO4. Are these soluble? NaCl is (by the same rule we just quoted), but what about SrSO4? Compounds of the sulfate ion are generally soluble, but Sr2+ is an exception: we expect it to be insoluble—a precipitate. Therefore, we expect a reaction to occur, and the balanced chemical equation would be:

Na2SO4(aq) + SrCl2(aq) → 2NaCl(aq) + SrSO4(s)

You would expect to see a visual change corresponding to SrSO4 precipitating out of solution (Figure 4.2 “Double-Replacement Reactions”).

A beaker full of a blue liquid. Some white solids are at the bottom.
Figure 4.2 “Double-Replacement Reactions.” Some double-replacement reactions are obvious because you can see a solid precipitate coming out of the solution.

Example 4.5

Problems

Will a double-replacement reaction occur? If so, identify the products.

  1. Ca(NO3)2 + KBr → ?
  2. NaOH + FeCl2 → ?

Solutions

  1. According to the solubility rules, both Ca(NO3)2 and KBr are soluble. Now we consider what the double-replacement products would be by switching the cations (or the anions)—namely, CaBr2 and KNO3. However, the solubility rules predict that these two substances would also be soluble, so no precipitate would form. Thus, we predict no reaction in this case.
  2. According to the solubility rules, both NaOH and FeCl2 are expected to be soluble. If we assume that a double-replacement reaction may occur, we need to consider the possible products, which would be NaCl and Fe(OH)2. NaCl is soluble, but, according to the solubility rules, Fe(OH)2 is not. Therefore, a reaction would occur, and Fe(OH)2(s) would precipitate out of solution. The balanced chemical equation is: 2NaOH(aq) + FeCl2(aq) → 2NaCl(aq) + Fe(OH)2(s)

Test Yourself

Will a double-replacement equation occur? If so, identify the products.

Sr(NO3)2 + KCl → ?

Answer

No reaction; all possible products are soluble.

Key Takeaways

  • A single-replacement reaction replaces one element for another in a compound.
  • The periodic table or an activity series can help predict whether single-replacement reactions occur.
  • A double-replacement reaction exchanges the cations (or the anions) of two ionic compounds.
  • A precipitation reaction is a double-replacement reaction in which one product is a solid precipitate.
  • Solubility rules are used to predict whether some double-replacement reactions will occur.

Exercises

Questions

  1. What are the general characteristics that help you recognize single-replacement reactions?
  2. What are the general characteristics that help you recognize double-replacement reactions?
  3. Assuming that each single-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Zn + Fe(NO3)2 → ?
    2. F2 + FeI3 → ?
  4. Assuming that each single-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Li + MgSO4 → ?
    2. NaBr + Cl2 → ?
  5. Assuming that each single-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Sn + H2SO4 → ?
    2. Al + NiBr2 → ?
  6. Assuming that each single-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Mg + HCl → ?
    2. HI + Br2 → ?
  7. Use the periodic table or the activity series to predict if each single-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. FeCl2 + Br2 → ?
    2. Fe(NO3)3 + Al → ?
  8. Use the periodic table or the activity series to predict if each single-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. Zn + Fe3(PO4)2 → ?
    2. Ag + HNO3 → ?
  9. Use the periodic table or the activity series to predict if each single-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. NaI + Cl2 → ?
    2. AgCl + Au → ?
  10. Use the periodic table or the activity series to predict if each single-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. Pt + H3PO4 → ?
    2. Li + H2O → ? (Hint: treat H2O as if it were composed of H+ and OH− ions.)
  11. Assuming that each double-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Zn(NO3)2 + NaOH → ?
    2. HCl + Na2S → ?
  12. Assuming that each double-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Ca(C2H3O2)2 + HNO3 → ?
    2. Na2CO3 + Sr(NO2)2 → ?
  13. Assuming that each double-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Pb(NO3)2 + KBr → ?
    2. K2O + MgCO3 → ?
  14. Assuming that each double-replacement reaction occurs, predict the products and write each balanced chemical equation.
    1. Sn(OH)2 + FeBr3 → ?
    2. CsNO3 + KCl → ?
  15. Use the solubility rules to predict if each double-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. Pb(NO3)2 + KBr → ?
    2. K2O + Na2CO3 → ?
  16. Use the solubility rules to predict if each double-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. Na2CO3 + Sr(NO2)2 → ?
    2. (NH4)2SO4 + Ba(NO3)2 → ?
  17. Use the solubility rules to predict if each double-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. K3PO4 + SrCl2 → ?
    2. NaOH + MgCl2 → ?
  18. Use the solubility rules to predict if each double-replacement reaction will occur and, if so, write a balanced chemical equation.
    1. KC2H3O2 + Li2CO3 → ?
    2. KOH + AgNO3 → ?

Answers

  1. One element replaces another element in a compound.
    1. Zn + Fe(NO3)2 → Zn(NO3)2 + Fe
    2. 3F2 + 2FeI3 → 3I2 + 2FeF3
    1. Sn + H2SO4 → SnSO4 + H2
    2. 2Al + 3NiBr2 → 2AlBr3 + 3Ni
    1. No reaction occurs.
    2. Fe(NO3)3 + Al → Al(NO3)3 + Fe
    1. 2NaI + Cl2 → 2NaCl + I2
    2. No reaction occurs.
    1. Zn(NO3)2 + 2NaOH → Zn(OH)2 + 2NaNO3
    2. 2HCl + Na2S → 2NaCl + H2S
    1. Pb(NO3)2 + 2KBr → PbBr2 + 2KNO3
    2. K2O + MgCO3 → K2CO3 + MgO
    1. Pb(NO3)2 + 2KBr → PbBr2(s) + 2KNO3
    2. No reaction occurs.
    1. 2K3PO4 + 3SrCl2 → Sr3(PO4)2(s) + 6KCl
    2. 2NaOH + MgCl2 → 2NaCl + Mg(OH)2(s)

Media Attributions

Figure 4.1

  • “Halogens on the Periodic Table” by David W. Ball © CC BY-NC-SA (Attribution NonCommercial ShareAlike)

Figure 4.2

  • “Copper solution” by Choij © Public Domain

Annotate

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Ionic Equations: A Closer Look
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Copyright © 2014

                                by Jessie A. Key

            Introductory Chemistry - 1st Canadian Edition by Jessie A. Key is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.
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