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Chapter Chemical Reactions

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					Chemical Reactions

    Reactions and Equations
     Chapter 10, Section 1
Evidence of Chemical
Reactions
   Chemical Reaction: the process by
    which the atoms of one or more
    substances are rearranged to form
    different substances
       Break down food to produce energy
       Produce natural fibers, cotton and wool
Representing Chemical
Reactions
   Reactants: the starting substances
   Products: the substances formed
    during the reaction

      reactant 1 + reactant 2  product 1 + product 2
    Reactants written to the                   Products are written to
                               Arrow is read
    left of the arrow                          the right of the arrow
                               as “react to
                               produce” or
                               “yields”.
   Symbols Used in Equations

Symbol                    Meaning
  +      Separates 2 or more reactants or products
        Separates reactants from products
  (s)    Identifies a solid state
  (l)    Identifies a liquid state
 (g)     Identifies a gaseous state
 (aq)    Identifies water solution
Example
   How would you write the equation that
    describes the reactions between carbon
    and sulfur to form carbon disulfide?
Example
   How would you write the equation that
    describes the reactions between carbon
    and sulfur to form carbon disulfide?
First write the chemical formulas for the reactants to
   the left of the arrow including their physical states.

                       C(s) + S(s) 
Example
   How would you write the equation that
    describes the reactions between carbon
    and sulfur to form carbon disulfide?
Finally write the chemical formula for the
  product, liquid carbon disulfide to the right of
  the arrow, indicating its physical state.

               C(s) + S(s)  CS2(l)
Practice Problems
   Write skeleton equations for the
    following word equations.
       Hydrogen(g) + bromine(g)  hydrogen bromide(g)
       Carbon monoxide(g) + oxygen(g)  carbon dioxide(g)
       Potassium chlorate(s)  potassium chloride(s) + oxygen(g)
Balancing Chemical Equations
   Chemical Equation: a statement that uses
    chemical formulas to show the identities and
    relative amounts of the substances involved
    in a chemical reaction.
Steps for Balancing Equations
1. Write the skeleton equation for the reaction.
     ex. H2(g) + Cl2(g)  HCl(g)
2. Count the atoms of the elements in the reactants.
     ex. H2  2 atoms of H
          Cl2  2 atoms of Cl
3. Count the atoms of the elements in the products.
     ex. HCl  1 atom H + 1 atom Cl
4. Change the coefficients to make the number of atoms of each
     element equal on both sides of the equation. Never change
     a subscript!!
     ex. H2(g) + Cl2(g)  2HCl(g)
Steps for Balancing Equations
5. Write the coefficient in their lowest possible ratio.
     ex. (2:2)  (1:1)
6. Check your work.
Example Problem
   Write the balanced chemical equation
    for the reaction in which sodium
    hydroxide and calcium bromide react to
    produce solid calcium hydroxide and
    sodium bromide. The reaction occurs in
    water.
Example Problem
   Write the balanced chemical equation
    for the reaction in which sodium
    hydroxide and calcium bromide react to
    produce solid calcium hydroxide and
    sodium bromide. The reaction occurs in
    water.

2NaOH(aq) + CaBr2(aq)  Ca(OH)2(s) + 2NaBr(aq)
Practice Problems
   Write chemical equations for each of the
    following reactions.
       In water, iron(III) chloride reacts with sodium
        hydroxide, producing solid Iron (III) hydroxide
        and sodium chloride.
       Liquid carbon disulfide reacts with oxygen gas,
        producing caron dioxide gas and sulfur dioxide
        gas.
       Solid zinc and aqueous hydrogen sulfate react to
        produce hydrogen gas and aqueous zinc sulfide.
Chemical Reactions

  Classifying Chemical Reactions
      Chapter 10, Section 2
Types of Chemical Reactions
   Synthesis
   Combustion
   Decomposition
   Single-replacement
   Double-replacement
Synthesis Reactions
   A chemical reaction in which two or
    more substances react to produce a
    single product.
       A + B  AB
           Ex. 2Na + Cl2  2NaCl
           Ex. CaO + H2O  Ca(OH)2
           Ex. 2SO2 + O2  2SO3
Combustion Reactions
   A reaction during which oxygen
    combines with a substance and
    releases energy in the form of heat and
    light.
       Ex. 2H2 + O2  2H2O
       Ex. C + O2  CO2
       Ex. CH4 + 2O2  CO2 + 2H2O
Practice Problems
   Write chemical equations for the following
    reactions. Classify each reaction into as
    many categories as possible.
       The solids aluminum and sulfur react to produce
        aluminum sulfide
       Water and dinitrogen pentoxide gas react to
        produce aqueous hydrogen nitrate.
       The gases nitrogen dioxide and oxygen react to
        produce dinitrogen pentoxide gas.
       Ethane gas (C2H6) burns in air, producing carbon
        dioxide gas and water vapor.
Decomposition Reactions
   A reaction in which a single compound
    breaks down into two or more
    elements or new compounds
       Ex. AB  A + B
           Ex. NH4NO3  N2O + 2H2O
           Ex. 2NaN3  2Na + 3N2
Practice Problems
   Write chemical equations for the
    following decomposition reactions.
       Aluminum oxide decomposes when
        electricity is passed through it.
       Nickel (II) hydroxide decomposes to
        produce nickel (II) oxide and water.
       Heating sodium hydrogen carbonate
        produces sodium carbonate, carbon
        dioxide and water.
Replacement Reactions –
Single & Double
   Single-replacement reaction: A
    reaction in which the atoms of one
    element replace the atoms of another
    element in a compound.
       Ex. A + BX  AX + B
           Ex. Cu + 2AgNO3  2Ag + Cu(NO3)2
     Replacement Reactions


   ** A metal will not
    always replace
    another metal
   Reactivity
    determines whether
    or not a metal will
    replace another
    metal.
   The most active
    metals, do not
    replace metals
Example Problem
   Predict the products that will result
    when these reactants combine and
    write a balanced chemical equation for
    each reaction.
       Fe + CuSO4
       Br2 + MgCl2
       Mg + AlCl3
Practice Problems
   Predict if the following single-
    replacement reactions will occur. If a
    reaction occurs, write a balanced
    equation for the reaction.
       2K + ZnCl2
       Cl2 + 2HF
       Fe + Na3PO4
Double Replacement
Reactions
   A reaction involving the exchange of
    positive ions between two compounds
    dissolved in water most often
    producing a precipitate, water or a gas
    (H2S, HCN, and CO2)
       Ex. AX + BY  AY + BX
           Ex. Ca(OH)2 + 2HCl  CaCl2 + 2H2O
           Ex. 2NaOH + CuCl2  2NaCl + Cu(OH)2
  Guidelines for Double-
  Replacement Reactions
Step                                 Example

1. Write the components of the       Al(NO3)3 + H2SO4
reactants in a skeleton equation.
2. Identify the cations and anions   Al(NO3)3 has Al3+ and NO3-
in each compound.                    H2SO4 has H+ and SO42-
3. Pair up each cation with the      Al3+ pairs with SO42-
anion from the other compound.       H+ pairs with NO3-
4. Write the formulas for the        Al2(SO4)3
products using the pairs from step   HNO3
3
5. Write the complete equation for   Al(NO3)3 + H2SO4  Al2(SO4)3 + HNO3
the double-replacement reaction.
6. Balance the equation.             2Al(NO3)3 + 3H2SO4  Al2(SO4)3 + 6HNO3
Practice Problems
   Write the balanced chemical equations for
    the following double-replacement reactions.
       Aqueous lithium iodide and aqueous silver nitrate
        react to produce solid silver iodide and aqueous
        lithium nitrate.
       Aqueous barium chloride and aqueous potassium
        carbonate react to produce solid barium
        carbonate and aqueous potassium chloride.
       Aqueous sodium oxalate and aqueous lead(II)
        nitrate react to produce solid lead(II) oxalate and
        aqueous sodium nitrate.
         Predicting Products of
         Chemical Reactions
Class of Reaction    Reactants                    Probable Products
Synthesis            Two or more substances       One compound
Combustion           A metal and oxygen           The oxide of the metal
                     A nonmetal and oxygen        The oxide of the nonmetal
                     A compound and oxygen        Two or more oxides
Decomposition        One compound                 Two or more elements
                                                  and/or compounds


Single-replacement   A metal and a compound       A new compound and the
                     A nonmetal and a compound    replaced metal.
                                                  A new compound and the
                                                  replaced nonmetal.
Double-              Two compounds dissolved in   Two different compound,
replacement          water                        one of which is often a solid,
                                                  water or a gas.
Chemical Reactions

  Reactions in Aqueous Solutions
      Chapter 10, Section 3
Aqueous Solutions
   Solutes: a substance dissolved in a
    solution
   Solvent: the substance that dissolves a
    solute to form a solution
   Aqueous solution: a solution in which
    the solvent is water
Reactions that Form
Precipitates
2NaOH(aq) + CuCl2(aq)  2NaCl(aq) + Cu(OH)2(s)
  In solution, these exist as ions. When their
   solutions are mixed, a precipitate of copper
               (II) hydroxide forms.
Ionic equations are written to show the details
                  of the reaction.
   2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq) 
          2Na+(aq) + 2Cl-(aq) + Cu(OH)2(s)
Reactions that From
Precipitates
   Complete ionic equation: an ionic
    equation that shows all of the particles
    in a solution as they realistically exist
   Spectator ions: ions that do not
    participate in a reaction
   Net ionic equations: ionic equations that
    include only the particles that
    participate in the reaction
Example Problem
   Write the chemical, complete ionic and
    net ionic equation for the reaction
    between aqueous solutions of barium
    nitrate and sodium carbonate that
    forms the precipitate barium carbonate.
Practice Problems
   Write chemical, complete ionic and net ionic equations for the
    following reactions that may produce precipitates. Use NR to
    indicate that no reaction occurs.
        Aqueous solutions of potassium iodide and silver nitrate are mixed,
         forming the precipitate silver iodide.
        Aqueous solutions of ammonium phosphate and sodium sulfate are
         mixed. No precipitate forms and no gas is produced.
        Aqueous solutions of aluminum chloride and sodium hydroxide are
         mixed, forming the precipitate aluminum hydroxide.
        Aqueous solutions of lithium sulfate and calcium nitrate are mixed,
         forming the precipitate calcium sulfate.
        Aqueous solutions of sodium carbonate and manganese (V)
         chloride are mixed, forming the precipitate manganese (V)
         carbonate.
Reactions That Form Water
   Double replacement reaction
   No evidence of chemical reaction is
    observable

      HBr(aq) + NaOH(aq)  H2O(l) + NaBr(aq)

               Complete Ionic Equation:
H+(aq) + Br-(aq) + Na+(aq) + OH-(aq)  H2O(l) + Na+(aq) + Br-(aq)
Reactions That Form Water
Crossing out the spectator ions leaves us with:
            H+(aq) + OH-(aq)  H2O(l)
              (net ionic equation)
Example Problem
   Write the chemical, complete ionic and
    net ionic equations for the reaction
    between hydrochloric acid and aqueous
    lithium hydroxide, which produces
    water.
Practice Problems
   Write chemical, complete ions and net ionic
    equations for the reactions between the
    following substances, which produce water.
       Sulfuric acid and aqueous potassium hydroxide
       Hydrochloric acid and aqueous calcium hydroxide
       Nitric acid and aqueous ammonium hydroxide
       Hydrosulfuric acid and aqueous calcium hydroxide
       Phosphoric acid and aqueous magnesium
        hydroxide
Reactions that Form Gases
   Double Replacement Reaction
   Commonly produce CO2, H2S and HCN

2HI(aq) + Li2S(aq)  H2S(g) + 2LiI(aq)

            Ionic Equation
2H+(aq) + 2I-(aq) + 2Li+(aq) + S2-(aq)
     H2S(g) + 2Li+(aq) + 2I-(aq)
Example Problem
   Write the chemical, complete ionic and
    net ionic equations for the reaction
    between hydrochloric acid and aqueous
    sodium sulfide which produces
    hydrogen sulfide gas.
Practice Problems
   Write chemical, complete ionic and net ionic
    equations for these reactions.
       Perchloric acid reacts with aqueous potassium
        carbonate
       Sulfuric acid reacts with aqueous sodium cyanide
       Hydrobromic acid reacts with aqueous ammonium
        carbonate
       Nitric acid reacts with aqueous potassium
        rubidium sulfide

				
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