AP Chem Kinetics
★ Chemical Kinetics
○ How fast a reaction occurs
○ Real life examples:
■ Airbags
■ Cooking
■ Rapid tylenol
■ Refrigerators
★ Collision Theory
○ For a reaction to occur, particles must collide
successfully
○ A successful collision has two components:
■ Particles must collide at the right geometry
● In order to break bonds and form
products
■ Particles must collide with enough energy
● In order to break the EA barrier
★ Effective Collision
○ A collision in which the particles meet with sufficient
energy to break activation energy barrier and an
orientation (geometry) that allows them to react
★ Four Factors that Create Effective Collisions
○ Concentration
■ Greater the number of molecules you have, the
greater amount of collisions will occur
○ Temperature
■ Higher temperature means particles collide
with greater kinetic energy, increasing the
rate of reaction
○ Surface Area
■ Increased surface area means greater
chances for collisions, and therefore faster
rates
○ Catalysts
■ Catalysts speed up reactions without being
used up
■ They lower the activation energy, meaning
more collisions to be successful
■ Lowers EA through the use of an alternate
pathway
★ What are the two viewpoints to measure rates of
reactions?
○ The increase in concentration of a product per unit
of time
○ The decrease in concentration of a reactant per
unit of time
★ Average Rate of Reaction
○ The change in the quantity of a reactant or product
that take place in a period of time
○
𝚫 𝚫𝚫𝚫𝚫𝚫𝚫𝚫𝚫𝚫𝚫𝚫𝚫𝚫
= 𝚫𝚫𝚫𝚫 𝚫𝚫 𝚫𝚫𝚫
𝚫 𝚫𝚫𝚫𝚫
○ Absolute value
★ Initial Rate of Reaction
○ Instantaneous rate @ t = 0
★ Instantaneous Rate
○ Rate @ a specific time
○ The slope of the straight line tangent to the curve at
a specific time
★ Rate Law
○ An equation that relates the rates of a reaction to
the concentration of reactants (and sometimes
catalysts) raised to various powers
○ EX: Rate = k [NO]2[H2]
★ Reaction Order
○ Order of a reactant species equals the exponent of
the concentration of that species in the rate law
★ Overall Order of Reaction
○ Sum of the exponents in a rate law
○ EX: Rate = k [NO]2[H2]
■ Overall order: third order
★ Determination of Rate Law
○ They must be determined experimentally and
cannot be obtained simply by looking at the
balanced equation
★ Rate of Reaction Stoichiometry
○ Multiply rate by mole ratio
○ EX: A + 2B ➞
3C + D
■ If the rate of disappearance is 4.06 x 10-4 for B,
what is the rate of appearance for C
4.06 x 10-
3
2
★ Key Info
○ Rates change over time:
■ Reactant rates decrease
■ Product rates increase
★ General Rate of Reaction
○ Take number/rate given and divide by the
coefficient
○ EX: 2A + 2B ➞
C + 2D
■ If the rate of disappearance for A is 2.2 x 10-4,
what is the
2.2 x
2
= 1.1 x
general rate?
★ Given a Rate Law, Determine Change in Rate with Change
in Concentration
○ EX: Rate = k[NO]2[H2]
■ If [NO] is tripled and [H2] is doubled, how will
the rate be affected?
● Plug in numbers into rate law
● Rate = [3]2[2]
● Rate = 18
● Rate increased by a factor of 18
★ Determining Rate Laws from Experimental Data
○ Sannies style!
■ Look at data chart and find 2 experiments
where the same reactant is “frozen”
■ Of the 2 experiments, find which has the large
rate and stack them on top of each other (large
on top)
■ Create a sannie
Rate (larger)
●
■ Cancel out the “frozen” reactant and the k and
solve for the exponent
■ Create another sannie for the other reactant
★ Finding k (rate constant)
○ Once rate law is found, plug in the numbers from any
experiment and solve for k
○ Units for k:
■ M1 - overall order t -1
■ “t” can be any unit of time
★ Rate Laws - What If? (experimental data)
○ After finding rate law from experimental data
○ EX: What if the concentration of A is 1.5 and the
concentration of B is 2.4? What is the rate?
■ Rate = k [A]2[B]
■ k = 2.76 x 10-4
■ Rate = 2.76 x 10-4 [1.5]2 [2.4]
★ Types of Rate Laws
○ Differential Rate Law:
■ Shows how the reaction rate changes with
concentration
■ Must be specific in how defined
○ Integrated Rate Law:
■ Shows how concentration changes with time
■ Shows how much reactant is left at a given time
■ Shows how long it will take
■ Graphical determination of the order
★ Integrated Rate Law
○ ln [A]t/[A]0 = -kt
■ Leftovers on top = [A]t
■ Initial conc on bottom = [A]0
■ k = rate constant
■ t = time ***must be in same units as k
■ ***remember the negative (-) in the formula
★ Uni-molecular
○ Single car accident
○1
○ EX: O3➞
O2 + O
★ Bi-molecular
○ Most common -two car collide/get in accident
○2
○ EX: HI + HI ➞
activated complex ➞
H2 + I2
★ Ter-molecular
○ Rare due to probability of orientation and energy
both being correct
○ 3 or more
★ Reaction Mechanism Rate Law Style
1. Cancel all intermediaries
■ Anything on both sides (exception for catalyst)
2. Total the overall reaction
3. Write rate law based on the slowest step
■ This is the ONLY time coefficients can be used
for a rate law
4. Intermediaries cannot be in a rate law (catalysts
can) so substitute using the reactants from
previous reaction into rate law
★ Reaction Mechanism - Catalysts
○ Catalyst is present in both the first step and final
step (beginning and end) b/c it’s not used up
○ Catalysts can be in the rate law
○ Reactant first, then a product
★ Reaction Mechanism - Slowest Step
○ Rate-determining step
○ If this is the first step, it is automatic rate law w/
coefficients
○ Has strongest bonds
○ Has strongest EA
★ Activation Energy (EA)
○ Minimum energy needed for a reaction to occur
○ Difference between the reactant and the top of the
hill (on a graph)
○ Affected (lowered) by the presence of a catalyst
★ What changes with TIME for a first-order reaction?
○ Rate of reaction
★ What does NOT change with TIME for a first-order
reaction?
○ Rate constant (k)
○ Half-life
★ Kinetics Graphing
○ Average rate from graph is just slope
○ Once a line goes flat, the reaction is complete
★ Temperature Reaction Rate
○ usually the higher the temperature the faster the
reaction rate, more molecules pass the Ea barrier
★ Activated Complex
○ a transitional structure that results from an
effective collision and that persists while old bonds
are breaking and new bonds are forming
○ Arrangement of atoms @ top of barrier/hill
★ Why does temperature increase reaction rate?
○ A higher fraction of molecules will have a kinetic
energy that is greater than the activation energy
★ Maxwell-Boltzmann Distribution
○ At higher temperatures, more molecules will the
adequate energy to react
○ This increases rate
○ Faster gas is wider graph and to the right
★ Find Order by Straight Line Plot
○ Find the straight line graph (usually out of 3)
○ Look at what is on the y-axis ([A], log or ln [A], 1/[A])
■ Zero Order:
● Slope = -k
● [A] just the concentration
■ First Order:
● Slope = -k
● Log or ln [A]
■ Second Order:
● Slope = k
● 1 / [A]
★ Half-life
○ Time it takes for half of it to disappear
○ t1/2 =
0.693
𝚫
★ Half-life question that asks about fraction of initial
○ EX: A reactant R is being consumed in a first-order
reaction. What fraction of the initial R is consumed
in 4.o half-lives?
100
Half-life
50
Half-life
25
Half-life
12.5
Half-life
6.25
3.125
Half-life