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12.5 Collision Theory
The half-lives for zero, first, and secondorder reactions are summarized in Table 12.2.
We should not be surprised that atoms must collide before they can react with each other.
Chemical bonds are formed when atoms are close together.
This simple premise is the basis for a very powerful theory that explains many observations regarding chemical kinetics.
The species must collide in an orientation that allows contact between the atoms that will become bonds in the product.
The collision must occur with enough energy to allow the electrons to rearrange and form new bonds.
Cars have catalysts that can be used to reduce pollutant.
Muzzle flash is a side reaction of gunpowder that can happen with many firearms.
The reaction occurs at high temperature and pressure if carbon monoxide and oxygen are present.
The second case is more likely to result in the formation of carbon dioxide, which has a central carbon atom bonding to two oxygen atoms.
The orientation of the collision is very important in creating the desired product of the reaction.
There are two collisions between carbon monoxide and oxygen.
The orientation of the colliding molecule affects whether a reaction will occur.
There is no guarantee that the reaction will form carbon dioxide if the collision takes place with the correct orientation.
Every reaction requires a certain amount of energy for it to move forward, yielding an activated complex along the way.
Figure 12.15 shows that even a collision with the correct orientation can fail to form a reaction product.
There is a possibility of carbon monoxide reacting with oxygen to form carbon dioxide.
Solid lines represent bonds, while dotted lines represent overlaps that may or may not become bonds as product is formed.
Carbon dioxide cannot form in the first two examples because the O double bond is not impacted.
If the third extra oxygen atom separates from the rest of the molecule, the third transition state will result in the formation of carbon dioxide.
It is not possible to identify a transition state or activated complex in most circumstances.
The gas-phase reaction is too rapid to separate any chemical compound.
As concentrations increase, collision theory explains why most reaction rates increase.
There are more molecules per unit of volume when the concentration of any reacting substance increases.
If the energy of the collisions is adequate, there will be a faster reaction rate.
The energy needed to form a product is provided by a collision of a reactant molecule with another reactant molecule.
The reaction will take a long time if the activation energy is larger than the average energy of the molecule.
The fraction of the molecule with the necessary energy will be large if the activation energy is less than the average.
The lost energy is transferred to other Molecules in order to reach the transition state.
The transition state is represented by the curve's peak.
The collision theory of reaction rates is accommodated in the Arrhenius equation.
The system doesn't have enough energy to overcome the barrier.
No reaction occurs in such cases.
The system has so much energy that every collision with the correct orientation can overcome the activation barrier and cause the reaction to proceed.
The reaction is almost instantaneous.
The Arrhenius equation describes a lot of what we have already discussed.
For two reactions at the same temperature, the slower reaction has a lower rate constant.
An increase in temperature has the same effect as a decrease in activation energy.
The Arrhenius equation shows H2 and I2.
We can simply pick two data entries using the experimental data presented here.
When a limited number of temperature- dependent rate constants are available for the reaction of interest, this method is very effective.
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