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Chapter 11 Chemical Reactions - A-Ordner

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Chapter 11 Chemical Reactions - A-Ordner Powered By Docstoc
					Chapter 11
Chemical Reactions
 Chemistry 2
Describing Chemical Reactions 11.1
 Writing Chemical Equations 11.1
• Reactants  products
  ▫  = yields, gives, or reacts to produce
• Iron + Oxygen  iron(III) oxide
• Hydrogen peroxide  water + oxygen
  ▫ Hydrogen peroxide decomposes to form water and oxygen gas
• Skeleton equation – chemical equation that does not indicate
  the relative amount of the reactant and products
• Fe(s) + O2(g)  Fe2O3(s)
  ▫   (s) = solid
  ▫   (l) = liquid
  ▫   (g) = gas
  ▫   (aq) = aqueous (dissolved in water)
• Catalyst – substance that speeds up the reaction but is not used
  up in the reaction
  ▫ Formula is written above the arrow
• Practice Problems Page 324 # 1-2
  Balancing Chemical Equations 11.1
• Coefficients – small whole #s placed in front of
  formulas to balance
• Each side of equation should have same # of each type
  of atom
• Law of Conservatio nof mass: mass is neither created
  or destroyed in chemical reaction
 ▫ Atoms rearranged
 ▫ Bonds broken and bond formed
          Here are some practice problems.

     1. __NaCl + __BeF2  __NaF + __BeCl2

  2. __FeCl3 + __Be3(PO4)2  __BeCl2 + __FePO4

   3. __AgNO3 + __LiOH  __AgOH + __LiNO3

        4. __CH4 + __O2  __CO2 + __H2O

      5. __Mg + __Mn2O3  __MgO + __Mn

Practice Problems: Page 327 # 3 – 4, Page 328 #5-6
Types of Chemical Reactions 11.2
         Classifying Reactions 11.2
• 5 types of reactions
  1. Combination or Synthesis – 2 or mores substances
     react to form a single new substance
    Group A metal + nonmetal = 2K + Cl2  2KCl
    2 nonmetals can have more than 1 product
      S + O2  SO2
      2S + 3O2  2SO3
    Transition metal and nonmetal can have more than 1
     product
      Fe + S  FeS
      2Fe + 3S  Fe2S3
    Practice Problems page 331 # 13 - 14
2. Decomposition Reactions – chemical change in which a
   single compound breaks down into two or more simpler
   products
 ▫   Opposite of synthesis
 ▫   2HgO  2Hg + O2
 ▫   Difficult to predict product
 ▫   Most required energy input (endothermic)
 ▫   Practice Problems page 332 # 15 – 16
3. Single-Replacement Reaction – chemical change in
   which one element replaces a second element in a
   compound
 ▫ 2K + 2H2O  2KOH + H2
 ▫ Both reactant and product contain of an element and a
   compound
 ▫ Activity series – list metals in order of decreasing reactivity =
   table 11.2 page 333
      Halogens = reactivity decreases down column
      Br2 + NaI  NaBr + I2
      Br2 + NaCl  no reaction
• Metals from Li to Na will replace H from acids and
  water
• Metals from Mg to Pb will replace H from acids only

• Practice Problems Page 334 # 17
4. Double Replacement Reactions – a chemical change
   involving an exchange of positive ions bvetween 2
   compounds
 ▫ To occur, one of the following is usually true:
    1 of produces is slightly soluble and precipitates from
     solution
      Na2S(aq) + CD(NO3)(aq)  CdS(s) + 2NaNO3(aq)
      ▫ CdS precipitated out
    One of products is a gas
    One product is a molecular compound
 ▫ Practice Problems Page 335 # 18 - 19
5. Combustion Reaction – chemical change in which an
   element or a compound reacts with oxygen often
   producing energy in the form of heat and light
 ▫ ALWAYS involves oxygen
 ▫ Often uses hydrocarbons
    Complete combustion form CO2 and water (and energy)
    Incomplete combustion forms CO
      Supply of oxygen is limited
 ▫ Practice Problems Page 337 # 20 - 21
  Predicting the Products of a Chemical
              Reaction 11.2
• # of elements/compounds reacting is a good indicator
  of possible reaction type
 ▫ Combination = 2 or more reactants  single product
 ▫ Decomposition = single compound  2 or more
   substances
 ▫ Single-replacement = element + compound  element +
   compound
 ▫ Double-Replacement = 2 ionic compounds  2 new
   compounds
 ▫ Combustion = Oxygen + hydrocarbon (usually)  water
   and carbon dioxide
Reactions in Aqueous Solution 11.3
           Net Ionic Equations 11.3
• Many of chemical reactions take place in water (aqueous
  solution)
  ▫ 70% Earth surface covered by water
  ▫ 66% of human body is water
• Complete Ionic equation – an equation that shows
  dissolved ionic compounds as dissociated free ions
  ▫ 2Na1+(aq) + SO42-(aq) + Ba2+(aq) + 2Cl1-(aq)  2Na1+(aq) +
    2Cl1-(aq) + BaSO4(s)
     Cross out ions that appear unchanged on both sides =
      spectator ions
  ▫ Na1+(aq) + SO42-(aq) + Ba2+(aq) + 2Cl1-(aq)  2Na1+(aq) +
    2Cl1-(aq) + BaSO4(s)
     Write the net ionic equation
  ▫ Ba2+(aq) + SO42-(aq)  BaSO4(s)
     Then balance
        Predicting the Formation of a
               Precipitate 11.3
• Mixing ionic compounds can sometimes form a
  precipitate (insoluble salt)
• Solubility Rules for Ionic Compounds
 ▫ Salts of alkali metals (1A) and ammonia (NH4)+= Soluble
 ▫ Nitrate (NO3)- salts and chlorate (ClO3)- salts = Soluble
 ▫ Sulfate (SO4)2- slats except compounds with Pb2+, Ag+,
   Hg22+, Ba2+, Sr2+, and Ca2+ = Soluble
 ▫ Chloride salts, except compound with Ag+, Pb2+, and
   Hg22+ = Soluble
 ▫ Carbonates (CO3)2-, Phosphates (PO4)3-,
   Chromates(CrO4)2-, Sulfides, and Hydroxides (OH)-
• Precipitate example:
 ▫ Na2Co3 (aq) + Ba(NO3)2(aq)  ?? Precipitate formed
    Separate = 2Na+(aq) + CO32- (aq) + Ba2+ (aq) + 2NO3-(aq)
      Would form NaNO3 and BaCO3
      ▫ Na = Alkali = soluble
      ▫ Nitrate salts = soluble Carbonates generally insoluble = BaCO3 will
        precipitate out
        ▫ Ba2+ (aq) + CO32- (aq)  BaCO3 (s)
• Practice Problems Page 343 # 28 - 29

				
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