Chapter 11 Chemical Reactions - Download as PowerPoint

							         Chapter 21
    “Chemical Reactions”

1
           All chemical reactions…
l        have two parts:
        1. Reactants = the substances you
           start with
        2. Products = the substances you
           end up with
l        The reactants will turn into the
         products.
l        Reactants  Products
    2
    Products
    Reactants




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              Chemical Reaction
    l   Here’s a simple       http://www.youtube.c
        reaction of water,      om/watch?v=AanR
        food coloring and       rqZx2Gk&NR=1&fe
        mineral turpentine.     ature=fvwp
        The author uses
        light so you can
        see the reaction.




4
         Antoine Lavoisier’s Contributions
    1770s - he experimented with mercury by placing a
      measured mass of solid mercury oxide into a
      sealed container. By adding heat, the red powder
      transferred into mercury metal, and a gas was
      produced. He found that no mass was lost:

        mercury (II) oxide       oxygen    plus   mercury
            10 g             =      .7 g   +       9.3 g



    The total starting mass of all reactants equals the
      total final mass of all products.



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               The Father of Modern Chemistry
    Lavoisier realized in Elements of
       Chemistry that substances needed
       to be named based on their
       composition. The International
       Union of Pure and Applied
       Chemistry (IUPAC) was formed in
       1919. This is the commission that
       coordinates guidelines for naming
       chemical compounds
       systematically. This is the
       commission that provides the name
       of an element before it receives a
       permanent name. Lavoisier also
       pioneered early experiments in
       biochemistry, medicine and sports
       science.




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             In a chemical reaction
l   Atoms aren’t created or destroyed (according
    to the Law of Conservation of Mass)
l   A reaction can be described several ways:
#1. In a sentence every item is a word
  Copper reacts with chlorine to form copper (II)
     chloride.
#2. In a word       equation some symbols used
        Copper + chlorine  copper (II) chloride
    7
             Symbols in equations?
l       the arrow (→) separates the reactants
        from the products (arrow points to products)
         –Read as: “reacts to form” or yields
l        The plus sign = “and”
l        (s) after the formula = solid: Fe(s)
l        (g) after the formula = gas: CO2(g)
l        (l) after the formula = liquid: H2O(l)
    8
       Symbols used in equations
l   (aq) after the formula = dissolved
    in water, an aqueous solution:
    NaCl(aq) is a salt water solution
l   used after a product indicates a
    gas has been produced: H2↑
l   used after a product indicates a
    solid has been produced: PbI2↓
    9
     Symbols used in equations
■         double arrow indicates a
 reversible reaction (more later)
     
■    ,     shows that
      
               heat
                   
 heat is supplied to the reaction
     Pt
■   is used to indicate a
 catalyst is supplied (in this case,
 platinum is the catalyst)
10
          What is a catalyst?
l     A substance that speeds up a
     reaction, without being
     changed or used up by the
     reaction.
l     Enzymes are biological or
     protein catalysts in your body.
    11
     Coefficients: Unit Managers
 Atoms are rearranged (never created or
   destroyed.) Numbers to the left of the
   formulas represent the number of units
   of each substance taking part in a
   reaction. Ex: 2NaOH is two units of
   sodium hydroxide.
 Read the analogy of making sandwiches
   on page 636 of your textbook.

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         Metals in the Atmosphere
     1) When oxygen reacting with
         iron breaks down the metal
         bond (causing rust.)



 2)      When oxygen reacts with
         aluminum, the resulting
         aluminum oxide adheres to
         and protects the metal.

 3)      When copper reacts with
         oxygen, the corrosion creates a
         blue-green coating called
         patina (i.e., the Statue of
         Liberty.)


13
          The Skeleton Equation
l    Uses formulas and symbols to
     describe a reaction
     –but doesn’t indicate how many;
      this means they are NOT
      balanced
l    All chemical equations are a
     description of the reaction.
    14
     Write a skeleton equation for:
1.   Solid iron (III) sulfide reacts with
     gaseous hydrogen chloride to form
     iron (III) chloride and hydrogen
     sulfide gas.
2.   Nitric acid (HNO) dissolved in water
     reacts with solid sodium carbonate
     to form liquid water and carbon
     dioxide gas and sodium nitrate
     dissolved in water.
15
     Now, read these equations:
Fe(s) + O2(g)  FeO2(s)

Cu(s) + AgNO3(aq)  Ag(s) + Cu(NO3)2(aq)

          Pt
NO2(g)   N2(g) + O2(g)



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     Balanced Chemical Equations
l   Atoms can’t be created or destroyed
    in an ordinary reaction:
     –All the atoms we start with we must
      end up with (meaning: balanced!)
l   A balanced equation has the same
    number of each element on both
    sides of the equation.
    17
          Rules for balancing:
1) Assemble the correct formulas for all the
   reactants and products, using “+” and “→”
2) Count the number of atoms of each type
   appearing on both sides
3) Balance the elements one at a time by
   adding coefficients (the numbers in front)
   where you need more - save balancing the
   H and O until LAST!
     (hint: I prefer to save O until the very last)
4) Double-Check to make sure it is balanced.
18
l    Never change a subscript to balance an
     equation (You can only change coefficients)
      – If you change the subscript (formula) you
        are describing a different chemical.
      – H2O is a different compound than H2O2
l    Never put a coefficient in the middle of a
     formula; they must go only in the front
          2NaCl is okay, but Na2Cl is not.
    19
     Practice Balancing Examples
l _AgNO3
  2          + _Cu  _Cu(NO3)2 + 2
                                 _Ag

l _Mg
  3       + _N2  _Mg3N2

l _P
  4      + _O2  _P4O10
           5


l _Na
  2       + _H2O  _H2 + _NaOH
            2            2


l    _CH4 + _O2  _CO2 + _H2O
            2            2

    20
        Types of Reactions
 Chemists have identified 5 main
    categories of chemical reactions:
 1) Combustion
 2) Synthesis
 3) Decomposition
 4) Single Displacement
 5) Double Displacement


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           Types of Reactions



     http://www.youtube.com/watch?v=tE4668aarck




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         #1-Combustion Reactions
l   Combustion means “add oxygen”
l   Normally, a compound composed of only
    C, H, (and maybe O) is reacted with
    oxygen – usually called “burning”
l   A combustion reaction occurs when a
    substance reacts with oxygen to produce
    energy in the form of heat and light.
l   Many combustion reactions can fit into
    other categories of reactions.
    23
     Combustion Reaction
 http://www.youtube.com/watch?v=Naej
   FHVjH2w




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       Synthesis Reactions
 lA very easy reaction in which 2 or
  more substances combine to form
  another substance:
           A + B -> AB
 Example: Water
       2H2 (g) + O2 (g) -> 2H2O (g)



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     #2 - Decomposition Reactions
l   decompose = fall apart
                  AB -> A + B
l   one reactant breaks apart into two or more
    elements or compounds.
             electricity
                
                  
l   NaCl          Na + Cl2
l   CaCO3    CaO + CO2
               
l   Note that energy (heat, sunlight, electricity,
    etc.) is usually required
    26
     Single Replacement Reactions
 When one element replaces another
  element in a compound:
 A + BC -> AC + B
 Ex:
     Cu (s) + 2AgNO3 (aq) -> Cu(NO3) 2 (aq) + 2Ag (s)
     A         B C       -> A C             + B




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     Single Replacement Reactions
 Single replacement can cause
   problems (ex: cooking spinach in an
   aluminum pan.)




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     Single Replacement Reactions
l    Metals will replace other metals (and they
     can also replace hydrogen)
l    K + AlN 
l    Zn + HCl 
l    Think of water as: HOH
      – Metals replace the first H, and then
        combines with the hydroxide (OH).
l    Na + HOH 
    29
         Single Replacement Reactions
l   We can even tell whether or not a single
    replacement reaction will happen:
     –Because some chemicals are more
      “active” than others
     –More active replaces less active
l   Activity Series of Metals
l   Higher on the list replaces those lower.

    30
       The “Activity Series” of Metals
Higher      Lithium
activity   Potassium   1) Metals can replace other
           Calcium
           Sodium
                          metals, provided they are
           Magnesium      above the metal they are
           Aluminum
           Zinc           trying to replace
           Chromium      (for example, zinc will replace lead)
           Iron
           Nickel      2) Metals above hydrogen can
           Lead
           Hydrogen       replace hydrogen in acids.
           Bismuth
           Copper
           Mercury     3) Metals from sodium upward
           Silver
Lower
           Platinum
                          can replace hydrogen in
activity
           Gold           water.
  31
     Synthesis, Decomposition and Single Replacement

 http://www.youtube.com/watch?v=dxl
   WtsFinTM




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     Double Replacement Reactions
l    Two things replace each other
            AB + CD -> AD + CB
     – Reactants must be two ionic
       compounds, in aqueous solution
         Ba(NO3) (aq) + K2SO4 (aq) -> BaSO4 (s) + 2KNO3 (aq)


l    NaOH + FeCl3 
     – The positive ions change place.
l    NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1
     = NaOH + FeCl3 Fe(OH)3 + NaCl
    33
#4 - Double Replacement Reactions
 l   Have certain “driving forces”, or reasons
      –Will only happen if one of the
       products:
      a) doesn’t dissolve in water and forms
       a solid (a “precipitate”), or
      b) is a gas that bubbles out, or
      c) is a molecular compound (which will
       usually be water).
 34
     Double Replacement Examples
 http://www.youtube.com/watch?v=opY
   3FLrPTa4




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       Complete and balance:
l   assume all of the following
    reactions actually take place:
     CaCl2 + NaOH 
    CuCl2 + K2S 
    KOH + Fe(NO3)3 
    (NH4)2SO4 + BaF2 

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         Practice Examples:
 l   H2 + O2 
 l   H2O 
 l   Zn + H2SO4 
 l   HgO 
 l   KBr + Cl2 
 l   AgNO3 + NaCl 
 l   Mg(OH)2 + H2SO3 
37
      Oxidation-Reduction Reactions
 A common characteristic to many chemical
   reactions is the tendency of the
   substances to lose or gain electrons.
 Oxidation – lost of electrons
 Reduction – gain of electrons
 These reactions often involve oxygen
   (highly reactive) as they pull electrons
   from metallic elements. Oxidation and
   reduction always works in pairs (called
   redox.)

38
       Chemical Reactions and Energy
 l   All chemical reactions release or
     absorb energy (thermal, light, sound,
     electricity.)
 l   Bonding takes energy. Bond
     formations release energy.
     Sometimes more energy is needed
     for one than the other.


39
         Exergonic Reactions
 l   In exergonic reactions, the chemical
     reactions release energy. Less
     energy is required to break the
     original bonds than is released when
     new bonds form. Example: glow
     sticks Heat is another example of
     exergonic reactions (ex: heat packs
     for sore muscles.)

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        Exothermic Reactions
 You should remember from studying states of
     matter that, when thermal energy is released, it is
     called an exothermic reaction. For example,
     burning wood (even rusting is exothermic, even
     though it is difficult to detect a temperature
     change.)
 It is exothermic reactions, through the burning of
     fossil fuels, that provide power used in most
     homes and industries. Unfortunately impurities
     (sulfur) can create pollution in our atmosphere.



41
         Endergonic Reactions
 l   An endergonic reaction occurs when the
     chemical reaction requires more energy to
     break bonds then to form new ones. The
     energy absorbed can be in the form of
     heat, light or electricity.
 l   Electricity is often used to supply energy
     to endergonic reactions. The
     electroplating of metals is an example:



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          Electroplating
 http://www.youtube.com/watch?v=1zel
   5wamQe0&feature=related




43
     Endothermic Reaction
 When energy is needed in the form of
  thermal energy, the reaction is
  endothermic. This can be either
  chemical or physical (adding Epsom
  salts to water absorbs energy and
  lowers the temperature of the water.)
  Cold packs are an example of an
  endothermic reaction.

44
             Catalysts
 You have already learned that a
  catalyst can be used to speed up a
  chemical reaction without being
  permanently changed (the mass of
  the product is the same, just the
  reaction is faster.) Catalyst can
  sometimes be recovered and reused.
  Example: Polymers used to make
  plastics and fibers.

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              Inhibitors
 While catalysts speed up reactions,
  inhibitors are used to prevent certain
  reactions from occurring. Example:
  food preservatives are added to
  prevent the chemical breakdown
  (spoilage) of certain foods.
 Catalysts and Inhibitors can only
  increase or decrease the rate of
  reaction.

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     Rate of Reaction Factors
 Before adding a catalyst or inhibitor,
  scientists must take into account
  other factors that may affect reaction
  rate:
  - Concentration of reactants
  - Pressure
  - Temperature


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