Reaction Rates and Equilibrium

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Reaction Rates and Equilibrium
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What is meant by the rate of a chemical reaction?

• Can also be explained as the speed of he
  reaction, it is the amount of time required for a
  chemical reaction to come to completion.
• Different reactions take different times
• Burning
• Aging
• Ripening
• Rusting

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Collision Theory

• Atoms, ions and molecules can react to form
  products when they collide provided that the
  particles are orientated correctly and have
  enough kinetic energy.
• The relative orientations of the molecules during
  their collisions determine whether the atoms are
  suitably positioned to form new bonds.
• Particles lacking necessary kinetic energy to
  react bounce apart when they collide.
• Imagine clay balls
• The minimum amount of energy that particles must
  have in order to react is called the activation
  energy

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Activated Complex
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Activation Energy

• Swedish Chemist Svante Arrhenius explored
  activation energy, Ea
• He found that most reaction rate data obeyed an
  equation based on three factors
• The fraction of molecules possessing the
  necessary activation energy or greater
• The number of collisions occurring per second
• The fraction of collisions that have the
  appropriate orientation
• Arrhenius equation
• k Ae-Ea/RT

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Factors that Affect Reaction Rates

• Temperature (Alkaseltzer activity)
• Raise temperature, faster reaction rate
• Lower temperature, slower reaction rate
• Higher temperatures make molecules move faster
  because they have more kinetic energy so reaction
  is more likely
• Particle Size/Surface Area (lycopodium demo)
• The smaller the particle size the larger the
  surface area.
• An increase in surface area, increases the amount
  of reactant exposed, which increases the
  collision frequency.
• Can be dangerous coal powder or gas particles
  reacting to our lungs.

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• Catalyst (demo)
• A substance that increases the rate of a
  reaction without being used up itself during the
  reaction.
• They help reactions to proceed at a lower energy
  than is normally required.
• Enzymes catalyze reactions in our body
• An inhibitor interferes with the action of a
  catalyst

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• Concentration
• The number of reacting particles in a given
  volume also affects the rate at which reactions
  occur.
• Cramming more particles into a fixed volume
  increases the concentration of reactants, the
  collision frequency and therefore, reaction rate.

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Rate Law Effect of Concentration on Rate

• One way of studying the effect of concentration
  on reaction rate is to determine the way in which
  the rate at the beginning of a reaction depends
  on the starting concentrations.
• a A b B ? c C d D
• Rate kAmBn
• The exponents m and n are called reaction orders
  and depend on how the concentration of that
  reactant affects the rate of reaction.

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Concentration and Rate

• If we compare Experiments 1 and 2, we see that
  when NH4 doubles, the initial rate doubles.

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Concentration and Rate

• Likewise, when we compare Experiments 5 and 6,
  we see that when NO2- doubles, the initial rate
  doubles.

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Concentration and Rate

• This means
• Rate ? NH4
• Rate ? NO2-
• Rate ? NH4 NO2-
• which, when written as an equation, becomes
• Rate k NH4 NO2-
• This equation is called the rate law, and k is
  the rate constant.

Therefore,
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Rate Laws

• A rate law shows the relationship between the
  reaction rate and the concentrations of
  reactants.
• The exponents tell the order of the reaction with
  respect to each reactant.
• Since the rate law is
• Rate k NH4 NO2-
• the reaction is
• First-order in NH4 and
• First-order in NO2-.

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Rate Laws

• Rate k NH4 NO2-
• The overall reaction order can be found by adding
  the exponents on the reactants in the rate law.
• This reaction is second-order overall.

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Reversible Reactions

• Until now, most of the reactions we have examined
  have gone completely to products. Some reactions
  are reversible meaning the reaction from
  reactants to products also goes from products to
  reactants at the same time.
• Consider the following reaction
• 2SO2 O2 2SO3
• In this reaction, SO2 and O2 are placed in a
  container. Initially, the forward reaction
  proceeds and SO3 is produced. The rate of the
  forward reaction is much greater than the rate of
  the reverse reaction. As SO3 builds up, it
  starts to decompose into SO2 and O2. The rate of
  the forward reaction is decreasing and the rate
  of the reverse reaction is increasing.
  Eventually, SO3 decomposes to SO2 and O2 as fast
  as SO2 and O2 combine to form SO3 (see Fig. 19.10
  from text). When this happens, the reaction has
  achieved chemical equilibrium.
• Chemical equilibrium is when the rate of the
  forward reaction the rate of the reverse
  reaction. It says nothing about the amounts of
  reactants and products at equilibrium. Simulation

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Equilibrium Constants

• When a system reaches equilibrium there is a
  mathematical relationship between the
  concentrations of the products and the
  concentrations of the reactants.
• aA bB cC dD
• Kc Cc Dd
• Aa Bb
• Ch 19 pg 545-548 Practice problems 8-13

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Factors Affecting Equilibrium Le Chateliers
Principle

• If a stress is applied to a system at dynamic
  equilibrium, the system changes to relieve the
  stress.
• Concentration
• Temperature
• Pressure

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Concentration

• Increasing the concentration of a substance
  causes the reaction to favor one reaction
  direction over the other to remove some of the
  extra substance.
• Decreasing the concentration of a substance
  causes the reaction to shift to produce some of
  the substance that was removed.
• For example 2SO2 O2 2SO3
• a) If SO2 increases by adding more, what
  happens the concentrations of all of the
  substances?
• SO2 goes up O2 goes down SO3 goes up.
• Adding SO2 causes that to go up. The system
  tries to get rid of that extra SO2 so some of the
  additional will react with O2 causing the O2 to
  decrease. This speeding up of the forward
  reaction causes more SO3 to be produced. Adding
  SO2 causes the forward reaction to be favored to
  use up excess SO2.so reaction shifts to the
  right.
• b) If SO3 decreases by removing some, what
  happens the concentrations of all of the
  substances?
• SO2 goes down O2 goes down SO3 goes
  down.
• Removing SO3 causes the reaction to favor the
  forward reaction to try to produce more so
  reaction also shifts to the right.

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Temperature

• Increasing the temp. causes the equilibrium
  position to shift in the direction that absorbs
  heat (favors the endothermic reaction).
  Decreasing the temp. causes the equilibrium
  position to shift in the direction that produces
  heat (favors the exothermic reaction).
• For example 2SO2 O2 2SO3 heat
• a) If temp increases it will cause the reaction
  to shift left (favor the reverse reaction) to try
  to remove the extra heat so
• SO2 goes up O2 goes up SO3 goes down.

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Pressure

• Affects systems containing gas particles since
  gasses are most affected by pressure. If
  pressure is increased, the reaction will shift to
  the side that contains the fewest number of gas
  particles and vice-versa (favor the reaction
  where most gas particles are reactants).
• For example 2SO2(g) O2 (g) 2SO3(g) heat
• a) If pressure is increased, the reaction will
  shift to the side with the fewest gas particles
  so it will shift to the right and
• SO2 goes down O2 goes down SO3 goes up.
• Chapter 19 pg 544 Practice problems 6,7