12.5 Collision Theory

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.