chem 12 1 11 1 13 ppt

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					  Catalysts



1.11- The effects of Cat. on Ea
1.12 The Effect of Cat. on RM
   1.13 Some uses of Cat.
                     Our Goals…
• compare the PE diagrams for a catalyzed and uncatalyzed
  reaction in terms of:
   –   reactants
   –   products
   –   activated complex
   –   reaction intermediates
   –   reaction mechanism
   –   ΔH
   –   activation energy
• identify platinum in automobile catalytic converters as a catalyst
• describe the effect of a catalyst on a number of reactions, such
  as:
   – decomposition of hydrogen peroxide (catalysts: manganese (IV) oxide,
     raw liver, raw potato)
   – the reaction of the oxalate ion with acidified potassium permanganate
     solution (catalyst: Mn2+)
   – the decomposition of bleach (catalyst: cobalt (II) chloride)
       CATALYSTS SPEED IT UP
• A catalyst is like adding a bit of magic to a
  reaction.
• Reactions need a certain amount of energy
  to happen. If they don't have it, oh well, the
  reaction probably can't happen.
• A catalyst lowers the amount of energy
  needed so that a reaction can happen
  easier.
• A catalyst is about energy; it doesn't have to
  be another molecule.
   – If you fill a room with hydrogen gas and oxygen
     gas, very little will happen. If you light a match in
     that room (or just a spark), all of the hydrogen and
     oxygen will combine to create water molecules. It
     is an explosive reaction.
• The energy needed to make a
  reaction happen is called the
  activation energy.
• As everything moves around,
  energy is needed.
• The energy a reaction needs is
  usually in the form of heat.
• When a catalyst is added,
  something special happens.
  – Maybe a molecule shifts it's
    structure.
  – Maybe that catalyst makes two
    molecules combine and they
    release a ton of energy. That
    extra energy might help another
    reaction to occur. In our earlier
    example, the spark added the
    activation energy.
      The Effects of Catalysts

• Earlier we defined a catalyst as a substance that
  speeds up a reaction. Let’s refine that definition a
  little…
• CATALYST: a substance which provides an overall
  reaction with an alternative mechanism having lower
  activation energy.
• Catalysts provide an alternative pathway by inserting
  different intermediate steps and lowering the activation
  energy for the reaction to occur. ∆H is not changed,
  but the “energy hump’ is lowered.
• Since the activation energy is lowered, a
  greater fraction of reactant molecules will have
  sufficient KE to form the activated complex.
  Since more reactant molecules can react in a
  certain time, the forward reaction rate
  increases.
                    Recall:
             REACTANTS ↔ PRODUCTS

• By lowering the “energy hump” for the forward
   reaction, we have also lowered it for the reverse
   reaction.
• Since a greater fraction of product molecules will have
   sufficient KE to form the activated complex, the
   reverse reaction rate increases also.
• Therefore:
If the forward reaction rate doubles, the reverse reaction
                       rate also doubles.
   Let’s examine a “real life” example
            utilizing catalysts
• The decomposition of formic acid has been
  extensively studied. At room temperature, the reaction
  is very slow with no noticeable activity. As soon we
  acidify the solution with sulfuric acid, the solution
  begins to bubble
                    What is happening?
• Some important things to remember about catalysts:
  – The catalyst is an active participant in a reaction which is
    regenerated in a later step of the reaction mechanism.
  – ∆H for the overall reaction is the same for both the catalyzed
    and un-catalyzed reaction; only the intermediate reactions
    differ.
  – All intermediates and catalysts cancel out when the
    individual steps are added up to get the overall reaction.
   Catalysis' and activation energy
                       MnO2 catalyzes
                       decomposition of
                         H2O2
                       2 H2O2 ---> 2 H2O + O2


Uncatalyzed reaction
Catalyzed reaction
          Uses of Catalysts
• Estimates are that 90% of all commercially
  produced chemical products involve
  catalysts at some stage in the process of
  their manufacture.
• In 2005, catalytic processes generated
  about $900 billion in products worldwide.
Energy processing
• Petroleum refining makes intensive use of catalysis for alkylation, catalytic cracking
  (breaking long-chain hydrocarbons into smaller pieces), naphtha reforming, steam
  reforming (conversion of hydrocarbons into synthesis gas). Even the exhaust from
  the burning of fossil fuels is treated via catalysis: Catalytic converters, typically
  composed of platinum and rhodium, break down some of the more harmful
  byproducts of automobile exhaust.
                                2 CO + 2 NO → 2 CO2 + N2
• With regards to synthetic fuels, an old but still important process is the Fischer-
  Tropsch synthesis of hydrocarbons from synthesis gas, which itself is processed via ,
  catalysed by iron. Biodiesel and related biofuels require processing via both
  inorganic and biocatalysts.
• Fuel cells ( Unit 5) rely on catalysts for both the anodic and cathodic reactions.

Bulk chemicals
• Some of the largest scale chemicals are produced via catalytic oxidation, often using
   oxygen. Examples include nitric acid (from ammonia), sulfuric acid (from sulfur
   dioxide to sulfur trioxide by the chamber process), terephthalic acid from p-xylene,
   and acrylonitrile from propane and ammonia.
• Many other chemical products are generated by large-scale reduction, often via
   hydrogenation. The largest-scale example is ammonia, which is prepared via the
   Haber process from nitrogen. Methanol is prepared from carbon monoxide.
• Bulk polymers derived from ethylene and propylene are often prepared via Ziegler-
   Natta catalysis. Polyesters, polyamides, and isocyanates via acid-base catalysis.
• Most carbonylation processes require metal catalysts, examples include the
   Monsanto acetic acid process and hydroformylation.
Food processing
• One of the most obvious applications of catalysis is the hydrogenation (reaction with
  hydrogen gas) of fats using nickel catalyst to give margarine. Many other foodstuffs
  are prepared via biocatalysis (see below).

Biology
• In nature, enzymes are catalysts in metabolism and catabolism. Most biocatalysts
   are protein-based, i.e. enzymes,
• Biocatalysts can be thought of as intermediate between homogenous and
   heterogeneous catalysts, although strictly speaking soluble enzymes are
   homogeneous catalysts and membrane-bound enzymes are heterogeneous. Several
   factors affect the activity of enzymes (and other catalysts) including temperature, pH,
   concentration of enzyme, substrate, and products. A particularly important reagent in
   enzymatic reactions is water, which is the product of many bond-forming reactions
   and a reactant in many bond-breaking processes.
• Enzymes are employed to prepare many commodity chemicals including high-
   fructose corn syrup and acrylamide.

In the environment
• Catalysis impacts the environment by increasing the efficiency of industrial
    processes, but catalysis also directly plays a direct role in the environment. A notable
    example is the catalytic role of Chlorine free radicals in the break down of ozone.
    These radicals are formed by the action of ultraviolet radiation on
    chlorofluorocarbons (CFCs).
                                    Cl· + O3 → ClO· + O2
                                    ClO· + O· → Cl· + O2
          Catalytic Converters
• Cars use this technology to reduce harmful emissions
  into the environment.
• NO(g) and NO2(g) emissions are changed to N2 (g) and
  O2 (g).
• CO(g) emissions are changed to less harmful CO2 (g).
              Video:-0
• http://www.youtube.com/watch?v=uGSLW
  ZU7YtE
        Practice Exercises
1. #56 – 61 page 34
2. #62-63 page 36

				
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posted:6/13/2012
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