Chemistry 140
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Chemistry 140. Chapter 5 Chemical Reactions. CHAPTER OUTLINE. 5.1 Chemical Equations 5.2 Types of Reactions 5.3 Redox Reactions 5.4 Decomposition Reactions 5.5 Combination Reactions 5.6 Replacement Reactions 5.7 Ionic Equations 5.8 Energy and Reactions

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Chemistry 140

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Chemistry 140

Chapter 5

Chemical Reactions


CHAPTER OUTLINE

5.1 Chemical Equations

5.2 Types of Reactions

5.3 Redox Reactions

5.4 Decomposition Reactions

5.5 Combination Reactions

5.6 Replacement Reactions

5.7 Ionic Equations

5.8 Energy and Reactions

5.9 The Mole and Chemical Equations

5.10 The Limiting Reactant

5.11 Reaction Yields


LEARNING OBJECTIVES/ASSESSMENT

When you have completed your study of this chapter, you should be able to:

1. Identify the reactants and products in written reaction equations, and balance the equations by inspection. (Section 5.1; Exercises 5.2 and 5.6)

2. Assign oxidation numbers to elements in chemical formulas, and identify the oxidizing and reducing agents in redox reactions. (Section 5.3; Exercises 5.10 and 5.15)

3. Classify reactions into the categories of redox or nonredox, then into the categories of decomposition, combination, single replacement, or double replacement. (Sections 5.4, 5.5, and 5.6; Exercise 5.20)

4. Write molecular equations in total ionic and net ionic forms. (Section 5.7; Exercise 5.30 a, b, & c)

5. Classify reactions as exothermic or endothermic. (Section 5.8; Exercise 5.34)

6. Use the mole concept to do calculations based on chemical reaction equations. (Section 5.9; Exercise 5.42)

7. Use the mole concept to do calculations based on the limiting‐reactant principle. (Section 5.10; Exercise 5.52)

8. Use the mole concept to do percentage‐yield calculations. (Section 5.11; Exercise 5.56)


Describing Chemical Reactions

  • chemical reaction - a change in matter in which a substance with new identities is formed

  • Indications of a chemical reaction

    • production of heat

    • production of light

    • change in color

    • production of a gas


Facts about chemical reactions.

  • Chemical reactions release or absorb energy.

    • release energy - exothermic

    • absorb energy - endothermic

  • Atoms are rearranged in a chemical change.

    • atoms cannot be created or destroyed

    • coefficients are adjusted to satisfy the law of conservation of matter

  • Particles must collide for a chemical reaction to occur.

    • reactant particles must come in contact with each other


How are reactions written?

  • 1. Equations represent facts - write a word equation.

  • 2. Substitute symbols and formulas for elements and compounds.

  • 3. Use charges of ions to write formulas for compounds.

  • 4. Balance the equation by adjusting coefficients.


Sample Equation 1

  • 1. hydrogen + oxygen water

  • 2. H2 + O2 H2O

  • 3. formulas are correct

  • 4. 2H2 + O2 2H2O


Sample Equation 2

  • 1. calcium carbonate  calcium oxide + carbon dioxide

  • 2.

  • 3.

  • 4


Practice Problems

  • 1. zinc + sulfur  zinc sulfide

  • Zn + S  ZnS

  • 2. sodium chloride + silver nitrate  silver chloride + sodium nitrate

  • NaCl +AgNO3  AgCl + NaNO3

  • 3. potassium + water  potassium hydroxide + hydrogen

  • 2K + 2H2O  2KOH + H2

  • 4. sodium hydroxide + hydrochloric acid (HCl)  sodium chloride + water

  • NaOH + HCl  NaCl + HOH


More Practice Problems

  • 5. magnesium bromide + chlorine  magnesium chloride + bromine

  • MgBr2 + Cl2 MgCl2 + Br2

  • 6. sodium chloride + sulfuric acid (H2SO4)  sodium sulfate + hydrogen chloride

  • 7. aluminum + iron(III) oxide  aluminum oxide + iron

  • 8. butane + oxygen  carbon dioxide + water


What information is in an equation?

  • qualitative - tells what is present

    • names, formulas, etc.

    • (s) solid, (l) liquid, (g) gas

    • (aq) aqueous - in water solution

  • quantitative - tells how much is present

    • coefficients represent quantity of moles

    • energy may be given

      • left side - endothermic

      • right side - exothermic


Describe the following equation in words.

  • C6H12O6(aq) + 6O2(g) 6CO2(g) + 6H2O(l) + 2870kJ

  • One mole of aqueous glucose reacts with 6 moles of oxygen gas to produce 6 moles of carbon dioxide gas, 6 moles of water and 2870kJ of energy is given for each mole of glucose reacting.


Types of Chemical Reactions

  • Combustion - oxygen combines with hydrocarbons

  • C2H6(g) + O2(g) CO2(g) + H2O(l)

  • Any substance that burns in air reacts with oxygen.


Synthesis

  • Synthesis - two or more substances combine to form more complex substances.

  • A + B  AB

  • Example: P4(s) + 6Cl2(g)  4PCl3(l)


Decomposition Reactions

  • Decomposition - complex substances are broken down into simpler substances.

  • AB  A + B

  • Example: 2HgO  2Hg + O2


Single Replacement Reactions

  • Single Replacement - a substance in one compound is replaced by another substance.

  • A + BC  AC + B

  • Example:

  • Mg(s) + 2HCl(aq)  MgCl2(s) + H2(g)


Double Replacement Reactions

  • Double Replacement - two substances exchange components.

  • AB + CD  AD + CB

  • Example:

  • NaCl + AgNO3  AgCl + NaNO3


Predicting Products of Decomposition Reactions

  • 1. Decomposition of a metallic carbonate yields a metallic oxide and carbon dioxide.

  • 2. Decompostion of a metallic chlorate yields a metallic chloride and oxygen.

  • 3. Decomposition of a metallic hydroxide yields a metallic oxide and water.

  • 4. Decomposition of a metallic nitrate yields a metallic nitrite and oxygen.


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