Chapter 8 - Chemical Reactions

Part 1 - Describing Chemical Change
 Part 2 - Classifying Reactions
 Part 3 - Predicting Products
Part 4 - Activity Series & Aqueous Reactions
 Part 5 - Endothermic & Exothermic Reactions
Return to Chemistry Page

Chapter 8 – Chemical Reactions
Part 1 – Notes: Describing Chemical Change
Objectives:    Identify, define, and explain: chemical equation, chemical statement, catalyst, coefficient, balanced equation, yield, reactant, product, chemical reaction, diatomic element, Law of Conservation of Mass.
                        Differentiate between chemical and algebraic equations and between chemical reactions and chemical equations.
                        List various common indicators of a chemical reaction.
                        Write equations describing chemical reactions using appropriate symbols.
                        Write and balance chemical equations given names or formulas of reactants and products in the reaction.
                        List elements that exist as diatomic molecules in their free state.
Text Reference:    Section 8.1 – pages 203-211

A chemical reaction is the process by which one or more substances change into new substance(s).
                Example:    solid calcium reacts with solid sulfur to form solid calcium sulfide.

A chemical equation is the written representation, in symbolic terms, of the overall change in a chemical reaction on an atomic level.
                Example:    Ca(S) +  S(S) -->  CaS(S)

Observations that indicate a chemical reaction may have occurred include (but are not limited to):
    


A reactant is (one of) the initial substance(s) in a chemical reaction; it reacts to form something new.  Reactants are always found on the left side of the equation.

A product is (one of) the resulting substance(s) in a chemical reaction; it is produced from the reactant(s).  Products are always found on the right side of the equation.  The product is always different from the reactant(s) and the properties of a product are different than the properties of the constituent reactant(s).

A “+” represents the combining of two substances in a chemical reaction.  It is commonly read plus, combines with, or reacts with.

A “-->” represents the split in the equation between reactants and products; it points in the direction of the final result.  It is commonly read  yields, forms, breaks down into, combines to form, or decomposes into.  It is NEVER read as “equal.”

(s), (l), (g), and (aq) indicate . . .

Adding heat to an equation is indicated with . . .

A catalyst is represented by . . .        .  A catalyst is . . .

Chemical equations are said to be balanced; this means there are an equal number of atoms of each type of element present on opposite sides of the yield sign.  Symbolic representations of chemical reactions need to be balanced in order to be a true chemical equation.  If a symbolic representation is not balanced, it is called a chemical statement or a skeleton equation.

Law of Conservation of Mass –


When balancing equations, you must make sure that MASS is conserved and that there are the SAME number of each type of atom on both sides of the yield.  Be careful not to confuse the subscripts and the coefficients when balancing equations.  NEVER change the subscript in a formula in an attempt to balance an equation.  Subscripts of a compound are there to balance the charges of the ions of indicate the correct number of atoms covalently bonded.  The correct compound cannot be changed to get an equation to balance.

KEEP IN MIND THAT WHEN YOU USE COEFFICIENTS TO BALANCE EQUATIONS, THEY MUST BE IN THE SMALLEST WHOLE NUMBER RATIO.

Note:    In this class, you will be given directions telling you to balance the chemical equations.  In actuality, you are balancing chemical statements and making them equations.  

Diatomic molecules:    two atoms of the same element that form molecules when they are in a free or uncombined state

Elements that form diatomic molecule in their free state:


Simple Rules to Follow for Basic Balancing:
            1.    Balance metallic elements first.
            2.    Balance nonmetallic elements next, except oxygen and hydrogen.
            3.    Balance hydrogen and oxygen last.

Writing and Balancing Chemical Equations

            1.    Convert from compound names to compound formulas.
                 - Use the correct subscripts to make the formula correct.  Once the compounds are correctly written, DO NOT change them!
            2.    Make sure the elements that form diatomic molecules are written with the subscript of two.
            3.    Balance the chemical statement with whole numbers coefficients in the smallest whole number ratio.
                 - Do NOT change the subscripts in the compound formulas.
                 - Simply add coefficients to illustrate the conservation of mass in the balanced chemical equation.


 
Chapter 8 – Chemical Reactions
Part 1 – Assignment: Describing Chemical Change
 Balance the following chemical equations using coefficients with the smallest whole number ratio.

1.    _____NH4NO2 --> _____H2O + _____N2

2.    _____H2 + _____N2 -->  _____NH3

3.    _____MgCO3 --> _____MgO + _____CO2

4.    _____P4 + _____Cl2 --> _____PCl5

5.    _____CrO3 --> _____Cr2O3 + _____O2

6.    _____IF5 + _____H2O --> _____HF + _____HIO3

7.    _____NH3 + _____O2 --> _____NO + _____H2O

8.    _____HgSO4 + ____Hg + ____NaCl --> ____Hg2Cl2 + ____Na2SO4    

9.    _____HBrO3 --> _____Br2O5 + _____H2O

10.    _____NO2 + _____H2O --> _____HNO3 + _____NO

11.    _____NH4NO3 --> _____N2O + _____H2O

12.    _____B2O3 + _____H2O --> _____H3BO3

Write and balance the following equations.

13.    calcium phosphate + aluminum sulfate --> calcium sulfate + aluminum phosphate

14.    potassium hydroxide -->  potassium oxide + water

15.    potassium iodide + lead (II) nitrate --> potassium nitrate + lead (II) iodide

16.    iron (III) oxide + carbon --> carbon monoxide + iron


To Top of Page     or     Return to Chemistry Page
Chapter8 – Chemical Reactions
Part 2 – Notes: Classifying Reactions
Objectives:    Identify, define, and explain: combination reaction, synthesis reaction, decomposition reaction, single replacement reaction, double replacement reaction, combustion reaction, and rapid oxidation.
                        Explain the importance of classifying reactions.
                        Identify and classify a reaction as synthesis, decomposition, single replacement, double replacement, or combustion.
Text Reference:    Section 8.2 (part) – pages 212-221

It is useful to classify chemical reactions because . . .


There are FIVE basic types of reactions that we will need to be able to classify.

A synthesis reaction is one in which a product is being created (or synthesized) from two or more elements.  It is also a reaction where a more complex compound is created from two or more simpler complexes.  There are always two or more reactants but only one product.  Synthesis reactions are also known as composition reactions or combination reactions.
                Example:    

                In other words:    


A decomposition reaction is one in which the single reactant is broken down into two or more elements or simpler compounds.  It is the reverse of a synthesis reaction.  It has only one reactant but two or more products.  In such a reaction, the single reactant is decomposed into its constituent parts.  
                Example:

                In other words:


A single replacement reaction is one in which an element that is a reactant replaces an element in the other reactant.  In other words, one partner is switched.  Key point: a resulting ionic compound must have a + and a – ion.  The switch cannot result in two positive ions or two negative ions forming a compound.
                Example:

                In other words:


A double replacement reaction (also known as a double displacement reaction or metathesis reaction) is one in which the positive ions of the two reactants switch places.  As products, each ion has a new and different partner that it originally had.
                Example:

                In other words:


                There are 3 types of products that drive double replacements reactions.  They are:


A combustion reaction is a rapid oxidation that usually produced a flame.  For our purposes (but not always), oxygen (from the air) is a reactant.  Hydrocarbons are compounds that contain hydrogen and carbon.  (Combustion reactions may also include hydrocarbons that also contain oxygen.)  When hydrocarbons combust (react with oxygen), the products are always carbon dioxide and water vapor.
                Example:

                In other words:




Chapter 8 – Chemical Reactions
Part 2 – Assignment: Classifying Reactions
For each of the following, write a balanced chemical equation.  Then, classify the reaction as SYN, DEC, SR, DR, or COMB.

_________    1.    iron + sulfur --> iron (II) sulfide

_________    2.    zinc + copper (II) sulfate --> zinc sulfate + copper

_________    3.    silver nitrate + sodium bromate --> silver bromate + sodium nitrate

_________    4.    potassium chlorate --> potassium chloride + oxygen

_________    5.    mercury (II) oxide --> mercury + oxygen

_________    6.    aluminum sulfate + sodium bicarbonate --> aluminum hydroxide + sodium sulfate + carbon dioxide

_________    7.    magnesium nitride + water --> magnesium hydroxide + ammonia (NH3)

_________    8.    copper (II) oxide + ammonia (NH3) --> copper + water + nitrogen

_________    9.    iron + water --> iron (III) oxide + hydrogen

_________    10.    iron (III) chloride + potassium hydroxide --> potassium chloride + iron (III) hydroxide

_________    11.    aluminum + hydrogen sulfate --> aluminum sulfate + hydrogen

_________    12.    aluminum sulfate + calcium hydroxide --> aluminum hydroxide + calcium sulfate


13.    What is the purpose of a catalyst?


14.    Explain why it is useful to classify reaction by their type.


15.    How do you predict the correct formula for the combination reaction between a nonmetal and a Group 2 metal?


16.    What is characteristic for every synthesis reaction?
To Top of Page     or     Return to Chemistry Page
Chapter 8 – Chemical Reactions
Part 3 – Notes: Predicting Products of Reactions
Objectives:    Predict the products of synthesis, decomposition, single replacement, double replacement, and combustion reactions.
                         List three types of products that drive double replacement reactions.
Text Reference:    Section 8.2  pages – (212-221) and 222-224

Since you are able to classify a reaction, you will use the patterns from the classifications to predict the products of the reaction.

Combustion Reactions:
        The products of the combustion of a hydrocarbon are always carbon dioxide and water vapor.

            Example 1:    C6H6 +          O2  -->

            Example 2:    C2H5OH  +          O2  -->

Synthesis Reactions:
        Since a synthesis reaction is one where two or more elements (or simple compounds) combine to form a single compound, the resulting product must contain all elements/ions present in the reactants without any leftover or additional elements/ions.  The product is a single product with balanced charges and the equation must be able to be balanced with correct coefficients.

            Example 3:    phosphorus + oxygen -->  diphosphorus pentoxide    (Recall oxygen is diatomic.)


            Example 4:    sodium + chlorine -->


Decomposition Reactions:
        Since a decomposition reaction is one where a single reactant breaks into two or more elements (or simple compounds), the resulting products must contain all elements/ions present in the reactant with no leftover.  The rules are the same for a synthesis – since they are opposite reactants.
    
            Example 5:    copper (I) chloride -->                (Recall chlorine is diatomic.)


            Example 6:    water -->


Single Replacement Reactions:
        In a single replacement reaction, either the positive or negative ion of one reactant is replaced by the other reactant (an element that forms either a positive or negative ion).  Only positive ions may replace other positive ions and only negative ions may replace other negative ions.
  
             Example 7:    aluminum + barium chloride -->


            Example 8:    copper (II) chloride + fluoride -->


Double Replacement Reactions (Metathesis Reactions):
        In a double replacement reaction, each reactant ion gets a new partner; the positive ions switch thereby giving everyone a new partner.  There are two reactants, each compound is made of two ions (one positive and one negative).  When positive ions switch, each positive ion must end up with a negative ion (not one to which it had been previously attached).  The compounds are written so there is a zero net charge and then the equation is balanced.

            Example 9:    sodium chloride + aluminum oxide -->


            Example 10:    potassium permanganate + calcium phosphate -->


            What types of products drive double replacement reactions?



Chapter 3 – Chemical Reactions
Part 3 – Assignment: Predicting Products of Reactions
     For each of the following, classify the reaction, predict the products, and balance the equation.

1.    KNO3  -->

2.    Fe2O3  +          H2 -->

3.    Zn  +          O2 -->

4.    CuS  +          KF  -->

5.    Al  +          Cl2  -->

6.    Ba(ClO3)2  +          NaOH  -->

7.    Fe  +          O2  -->

8.    AuCl3  -->

9.    ZnI2  +          Br2  -->

10.    NH3  -->

For each of the following, write the reactants, predict the products, classify the reaction, and balance the equation.

11.    sodium oxide  +  iodine  -->


12.    calcium carbonate -->


13.    calcium sulfate  +  sodium nitrate -->


14.    hydrogen sulfite + calcium -->


15.    aluminum oxide + magnesium fluoride -->
To Top of Page     or     Return to Chemistry Page
Chapter 8 – Chemical Reactions
Part 4A – Notes: Predicting Products and the Activity Series
Objectives:    Define and explain activity series.
                         Use the activity series to predict is a reaction will occur.
Text Reference:    Section 8.2 – page 217

In order to correctly predict the products of a single replacement reaction, the chemical activity of the elements that will potentially switch places needs to be considered.  In order for an element to replace another, it must be ore active.  
A more active element will replace a less active element.  If this is not the case, then no reaction occurs.

Examine the activity series and note any patterns:
ACTIVITY SERIES
METALS
NONMETALS
lithium
 More
potassium
 Active
calcium
 V
sodium
 V
magnesium
 V
 fluorine
aluminum
 V
zinc
 V
chromium
 V
 chlorine
iron
 V
nickel
 Decreasing
tin
 Activity
 bromine
lead
 V
HYDROGEN
 V
copper
 V
 iodine
mercury
 V
silver
 V
platinum
 Less
gold
 Active


For example, for a set of reactants: magnesium and copper sulfate, you need to determine if the magnesium will replace the copper ion.  You check the activity series.  You find that the magnesium is more active than the copper and therefore the magnesium will replace the copper and the reaction will proceed.

Examples:    Will the following reactions occur?  If yes, what are the products?  If no, write NR.

1.    chromium + lead (II) chloride

2.    zinc + potassium hydroxide

3.    magnesium + sulfuric acid

4.    iodine + sodium chloride

5.    fluorine + sodium chloride



Chapter 8 – Chemical Reactions
Part 4B – Notes: Aqueous Reactions
Objectives:    Identify, define, and explain: complete ionic equation, spectator ion, net ionic equation, solubility rules, aqueous reaction, and precipitate.
                        Write and balance complete ionic equations and net ionic equations.
                        Use solubility rules to predict the precipitate formed in a double replacement aqueous reaction.
Text Reference:    Section 8.3 – pages 225-228

Deciding whether a double replacement reaction will occur naturally is actually a matter of predicting whether an insoluble product can form.  Solubility rules are used to determine this.  The solubility rules are a listing of the solubility of various substances in water.  Some compounds are listed as soluble and other compounds are listed as insoluble.  If one of the products formed is insoluble, it is this insoluble product that is the precipitate that forms in the reaction.  If no insoluble precipitate forms, then there is no aqueous reaction.
SOLUBILITY RULES TABLE


Examples:

1.    Mix silver nitrate and sodium bromide.  Does a reaction occur?

2.    Mix sulfuric acid and barium chloride.  Does a reaction occur?

3.    Mix lithium chloride and copper (II) acetate.  Does a reaction occur?



Chapter 8 – Chemical Reactions
Part 4 –Assignment: Writing and Balancing Reactions
 For each of the following, write the reactant(s), formula, predict and write the product formula, then classify and balance.

1.    potassium phosphate + magnesium chloride -->

2.    copper (II) nitrate + ammonium hydroxide -->

3.    lead (II) hydroxide -->

4.    ammonium sulfide + lead (II) nitrate -->

5.    sodium hydroxide + calcium nitrate -->

6.    water + diphosphorus pentoxide --> phosphoric acid

7.    chromium (II) chloride + sulfuric acid -->

8.    calcium chloride + chromium (III) nitrate -->

9.    iron (III) hydroxide -->

10.    potassium chlorate + lead (II) nitrate -->

11.    aluminum fluoride + sulfuric acid -->

12.    zinc arsenate + hydrochloric acid -->

13.    sodium carbonate -->

14.    copper (II) chlorate -->

15.    lithium + water -->

The following reactants are mixed in water.  Write and balance the following as complete and net ionic equations.

16.    copper (II) sulfate reacts with iron (III) chloride in an aqueous environment



17.    iron (III) sulfate reacts with ammonium sulfide in an aqueous environment



18.    potassium carbonate reacts with strontium nitrate in an aqueous environment



19.    sodium iodide reacts with silver nitrate in an aqueous environment



20.    aluminum sulfate reacts with calcium chloride in an aqueous environment
To Top of Page     or     Return to Chemistry Page
Chapter 8 – Chemical Reactions
Part 5 – Notes and Questions: Endothermic and Exothermic Reactions

Objectives:    Identify, define, and explain: endothermic reaction, exothermic reaction, and relative energy changes.
                        Explain the relative energies of reactants versus products of endothermic and exothermic reactions.
                        Classify a reaction as endothermic or exothermic.
                        Qualitatively describe the change to the system and the surroundings in an endothermic or exothermic reaction.
Text Reference:    Section 11.1 – pages 293-295

Reactions may be endothermic or exothermic.

1.    What is an endothermic reaction?


        How do you known when a chemical equation represents an endothermic reaction?


2.    What happens to the reaction system and the surroundings during an endothermic reaction?  (Respond in terms of heat.)


3.    Discuss the relative energies for reactants and products in an endothermic reaction.


4.    What is an exothermic reaction?


        How do you know when a chemical equation represents an exothermic reaction?


5.    What happens to the reaction system and the surroundings during an exothermic reaction?  (Respond in terms of heat.)


6.    Discuss the relative energies for reactants and products in an exothermic reaction.


7.    Based on energetic considerations only, which is a more probable reaction type: endothermic or exothermic?  Why?


8.    Classify the following as endothermic or exothermic.  Then balance the equation.

        a.    ____________________        HI  +  energy  -->          H2  +          I2

        b.    ____________________        H2O2  -->          H2O  +          O2  +  energy

        c.    ____________________        Na  +          H2O2  -->          NaOH  +          H2  +  energy

        d.    ____________________        Cu  +          O2  +  energy  -->          CuO

        e.    ____________________        NH3  -->          N2  +          H2  +  energy

        f.    ____________________        Mg  +          HCl  -->          MgCl2  +          H2  +  energy



Chapter 8 – Chemical Reactions
Part 5 – Assignment: Combustion Reactions
For each of the following sets of reactants, write the products and balance the equation.

1.    C6H6  +          O2  -->

2.    C12H22  +          O2  -->

3.    C7H13OH  +          O2 -->

4.    C12H22O11  +          O2 -->

5.    C8H17OH  +          O2  -->

6.    C6H12O6  +          O2  -->

7.    C9H20  +          O2  -->

8.    C11H22OH  +          O2  -->

9.    Text problem – page 234 #65.  Alkanes are hydrocarbon molecules that have the general formula CnH2n+2.  The graph shows the number of oxygen, carbon dioxide, and water molecules needed to balance the equations for the complete combustion of every alkane having from one to ten carbon atoms.

        Combustion of an alkane:    Alkane + ____O2  -->   ____CO2  +  ____H2O

        a.    Use the graph to write balanced equations for the combustion of C5H12 and C2H20.


        b.    Extrapolate the graph and write balanced equations for the combustion of C12H26 and C17H36.


        c.    The coefficient for O2 in the general equation is:    n + (square root [(n + 1)/2])
               What are the coefficients for CO2 and H2O, with respect to n?
To Top of Page     or     Return to Chemistry Page