Studied by 16 people
rate
A measure of change over time.
particle theory
All matter is in constant motion with the speed represented by the average kinetic energy (temperature) of the system.
kinetic energy
Energy associated with the motion of molecules.
collision theory
All particles in motion will randomly collide with one another. For a collision to result in a reaction, the collision must be effective. An effective collision (one that results in the formation of products) must satisfy two conditions.
effective collision
Reactant orientations must be favourable.
Collision must occur with sufficient energy.
Maxwell-Boltzmann Curve
In most reactions, only a small fraction of the total collisions have sufficient energy for a reaction to occur. The collision energy depends on the kinetic energy of the colliding particles.
Plotting the number of particles versus the kinetic energy of the particles gives a curve like the one below:
NOTE: The curve represents the distribution of the kinetic energy of collisions at a given temperature. The activation energy is independent of temperature (it does not change when temperature changes!).
potential energy diagram
Shows the relative potential energies of reactants, transition states and products as a reaction progresses.
potential energy
Energy within molecules (i.e. the bonding energy).
activation energy
The minimum amount of kinetic energy that the reactants must collide with to make the reaction proceed (energy barrier the reactants must overcome to form products).
Ea
activated complex/transition state
A temporary arrangement of atoms that form as bonds are breaking and new bonds are forming (highly unstable).
endothermic
exothermic
surface area
When increasing this (grinding, slicing, dissolving) of solid reactants, it exposes more molecules to each other than before, hence increasing the frequency/probability of effective collisions, increasing the rate (opposite occurs when this is small).
concentration
With an increase in the number of reactant molecules present, there is an increase in the frequency/probability of effective collisions occurring, hence increasing the rate (the reverse is true when this is low).
temperature
By increasing this, the molecules have more kinetic energy which increases the frequency/probability of effective collisions. With more energy, the activation barrier can be more easily overcome more often, increasing the rate (the opposite occurs when this is low).
volume/pressure
By decreasing this factor of a container and and increasing the other, the molecules have less space to move and are in closer proximity to one another, hence increasing the frequency/probability of effective collisions, increasing the rate (increasing the former and decreasing the latter decreases the rate).
catalyst
Allows for an alternate pathway of lower activation energy so that more collisions can overcome the barrier more often, increasing the rate. They are not consumed in reactions and can be reused.
mol/Ls
Overall order: 0
(write units of the rate constant k)
/s
Overall order: 1
(write units of the rate constant k)
L/mols
Overall order: 2
(write units of the rate constant k)
L^2/mol^2s
Overall order: 3
(write units of the rate constant k)
rate law equation
Relates the rate of a reaction to the concentrations of the reactants, each concentration is expressed with an order (exponent).
m and n are the reactant orders determined from experimental data
The rate constant is k, which makes the two sides of the equation equal to each other (units of k will change depending upon the overall order of the reaction)
reaction mechanism
The overall sequence of elementary steps by which a chemical reaction occurs.
A reaction that occurs in two or more elementary steps is called a multistep or complex reaction.
rate-determining step
When a reaction takes place in a series of steps, the overall rate of the chemical reaction is only as fast as the rate of its slowest step (the step with the largest activation energy).
The coefficient of each reactant in this is the order of each reactant in the rate law equation!
reverse endothermic
reverse exothermic
reaction intermediates
Stable chemical compounds that do NOT appear in the overall reaction.
rate-determining step
The overall rate of the chemical reaction is only as fast as the rate of its slowest step (the step with the largest activation energy). The rate law equation is dependent on this step; the coefficient of each reactant in the this step is the order of each reactant in the rate law equation!
Note 
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