Chemistry: Kinetics
Reaction co-ordinate: the measure of the extent to which the reaction was processed forward.
Activation energy: the minimum energy that must be available for collision to lead a reaction.
ΔH: net potential energy change
Transition state: the potential energy the reactant must reach before products can be formed.
Activated complex:(it is also known as the transition state) an unstable chemical species containing partially broken and partially formed bonds. Respenting the maximum potential energy point in the change.
Elementary steps: the smaller steps that a reaction is broken down into.
Reaction mechanism: a series of elementary steps that combine to makes up an overall reaction.
Rate-determining step: the slowest step in a reaction.
Relative rates:is proportional to the coefficient of the reactants
Assumptions of collision theory:
-the overall reaction probably doesn't describe the collision required for a reaction.
-bimolecule collision or unimolecule dissociation reactions are the most probable collision which can explain the rapid reactions observed. Therefore a reaction is most likely to be a serious of bimolecule or unimolecule reactions which add to give the overall reaction.
Order of reaction: the exponents in the rate raw describe the mathematical dependence of rate on initial concentration.
Overall order of reaction: the sum of order for each reactant.
Logarithms:
Factors that affect the reactions: concentration, catalyst, temperature, surface area
Rate law: the rate will always be proportional to the product of the initial concentration of the reactant, where these concentrations are raised to some exponential values. R=k[A]m[B]n
Intermediate: an unstable molecule formed in a reaction mechanism that has a very short life span.
Collision theory: the reaction occurs between two molecules if they collide at the correct orientation and the energy of the collision is sufficient to break the chemical bonds.
Reaction co-ordinate: the measure of the extent to which the reaction was processed forward.
Activation energy: the minimum energy that must be available for collision to lead a reaction.
ΔH: net potential energy change
Transition state: the potential energy the reactant must reach before products can be formed.
Activated complex:(it is also known as the transition state) an unstable chemical species containing partially broken and partially formed bonds. Respenting the maximum potential energy point in the change.
Elementary steps: the smaller steps that a reaction is broken down into.
Reaction mechanism: a series of elementary steps that combine to makes up an overall reaction.
Rate-determining step: the slowest step in a reaction.
Relative rates:is proportional to the coefficient of the reactants
Assumptions of collision theory:
-the overall reaction probably doesn't describe the collision required for a reaction.
-bimolecule collision or unimolecule dissociation reactions are the most probable collision which can explain the rapid reactions observed. Therefore a reaction is most likely to be a serious of bimolecule or unimolecule reactions which add to give the overall reaction.
Order of reaction: the exponents in the rate raw describe the mathematical dependence of rate on initial concentration.
Overall order of reaction: the sum of order for each reactant.
Logarithms:
Factors that affect the reactions: concentration, catalyst, temperature, surface area
Rate law: the rate will always be proportional to the product of the initial concentration of the reactant, where these concentrations are raised to some exponential values. R=k[A]m[B]n
Intermediate: an unstable molecule formed in a reaction mechanism that has a very short life span.
Collision theory: the reaction occurs between two molecules if they collide at the correct orientation and the energy of the collision is sufficient to break the chemical bonds.
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