Chemical Reactions

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Chemical Reactions

• How Fast? Why Use Catalysts?

2
Rate of Reaction

• Under a given set of conditions, each chemical
  reaction requires a characteristic amount of time
  to produce a given amount of product
• Rate of the reaction - the amount of a
  participant undergoing change per unit time, such
  as lbs/hour, tons/day, etc.

3
Collision Theory

• For reaction to occur, particles of the reactants
  must collide with sufficient energy to form a
  transition state in which all the reacting atoms
  are in a proper orientation
• Transition State - unstable state produced by the
  collision of reacting molecules - all reacting
  atoms are held together by weak bonds

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Transition State

• 2 HI(g) H2(g) I2(g) ?H 3 kcal
• Reactants Products
• 2 HI(g) H - H H2(g) I2(g) ?H 3
  kcal Reactants ?????? Products
  I - I
• Transition State

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Energy of Activation, Ea

• The minimum energy required to produce the
  transition state is called the energy of
  activation (Ea) for the reaction
• Following is a graph, called an energy profile,
  for the decomposition of HI into H2 and I2, where
  ?H 3 kcal and Ea 41 kcal

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Energy Profile of Decomposition of HI
2 HI(g) H2(g) I2(g) ?H 3 kcal and Ea
41 kcal
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Significance of Energy of Activation
As Ea increases, the fraction of collisions that
can produce products decreases. Therefore, the
higher the energy of activation,
the slower the rate of the reaction.
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Orientations of Transition State
Orientation 1 is a proper orientation - capable
of producing products
Orientation 2 is an improper orientation - not
capable of producing products
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Factors Which Influence the Rates of Chemical
Reactions

• 1. Nature of the Structure and Bonding of the
  Reactants
• a. Different reactions involving substances
  with different bonding require different energies
  of activation
• b. Different reactions involving substances
  with different structures will have different
  proper orientations

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1. Structure and Bonding of the Reactants

• C6H14 Br2 C6H13Br HBr (Slow) Hexane
• C6H12 Br2 C6H12Br2 (Fast) Hexene
• Difference in rates due to difference in bonding
  (changes Ea) and difference in structure (changes
  the geometry of the proper orientation)

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2. Surface Area

• An increase in the surface area of the more
  condensed reactant causes an increase in the rate
  of a heterogeneous reaction (a reaction involving
  reactants in more than one phase such as a
  gas and a solid)
• Lycopodium powder burns explosively when
  dispersed in air

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3. Concentrations of Reactants

• Increasing the concentration of a reactant
  frequently increases the rate of a reaction
• The Collision Theory explains this as being due
  to increased collision frequency

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4. Temperature

• Increasing temperature increases the rates of all
  reactions
• A 10 oC increase in temperature causes the rates
  of many reactions to double or triple

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Continued...

• Due, in part, to increased collision frequency as
  the molecules move faster at the higher
  temperature
• Due mostly to the increase in the fraction of
  molecules possessing the energy of activation

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Temperature and Energy of Activation
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5. Catalyst

• A catalyst is a substance whose presence alters
  the rate of a reaction without undergoing
  permanent change itself
• Adding KI(s) to 2 H2O2(aq) 2 H2O(l) O2(g)
  ?H -50 kcal increases rate, but all KI(s) can
  be recovered unchanged after reaction - KI is a
  catalyst

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Reaction Without Catalyst
2 H2O2(aq) 2 H2O(l) O2(g) ?H -50 kcal
High Ea since two large molecules must collide
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Reaction with Catalyst
H2O2 I- IO- H2O H2O2
IO- H2O I- O2 2 H2O2
2 H2O O2 Fast, Ea 25 kcal
Collision of peroxide and iodide has lower Ea,
faster rate.
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Heterogeneous Catalysis

• Heterogeneous catalyst is in a different state
  from the reactants usually a solid while
  reactants are gases or liquids
• Provides a surface to which reactants are
  attracted which, in turn, weakens bonds in
  reactants, lowers Ea, and increases rate of
  reaction
• MnO2(s) catalyzes peroxide reaction