Chapter 13: Chemical Kinetics Flashcards
chemical kinetics: study of … of chemical reactions and their …
rates of chemical reactions; reaction mechanisms
rate=
Δc/Δt
the reactant concentration decreases during the reaction, so Δ[reactant] is
negative
by convention, the rate of reaction is always expressed as a
positive number
the average rate of reaction is equal to the change in … divided by the …
concentration; time interval
as the interval between measurements becomes smaller, the average rate approaches the
instantaneous rate
when a rate is measured over a time interval, it is called an
average rate
instantaneous rate: the … to the curve at a particular time pt
slope of the tangent
the rate of reaction does not depend on
which species is measured
the rate of reaction is the absolute value of the rate of change of the concentration a substance divided by its ….
stoichiometric coefficient
the rate of a reaction is strongly influenced by the
concentrations of the reacting species
for the equation: aA + bB→products:
rate=
k[A]^x[B]^y
overall order:
sum of the individual orders
k is the
rate constant
initial rate method: repeat an experiment several times with different known … and evaluate how th reaction rate changes with …, important to measure the .. of the reaction
ratios of reactants; concentration; initial rate
the rate law cannot be predicted from the
reaction stoichiometry
rate laws can also be determined by expaning how the concentration of a reactant changes with … during the course of a single experiment
time
(zero-order rate laws) some reactions show rates that are …. of the concentration of reactants
independent
(zero-order rate laws) reaction rate=
k
(zero-order rate laws) the units of a zero-order rate constant are:
mol/L*s
(zero-order rate laws) if the graph of reactant concentration vs. time is a straight line, the reaction obeys
zero-order kinetics
(first-order rate laws) some reactions show rates that are …. to the concentration of the reactant
proportional
(first-order rate laws) reaction rate=
-Δ[A]/ Δ t → k[A]
(first-order rate laws) first-order rate constant:
s^-1
(first-order rate laws) if the graph of the … of reactant concentrations vs. time is a straight line, the reaction obeys first-order kinetics
natural log (ln)
(first-order rate laws) reaction rate= k[A] is referred to as the … of the rate law
differential form
(first-order rate laws) the rate law can be written in the following integrated form:
[A] =[A]0e-kt
(first-order rate laws) for [A] =[A]0e-kt,
A0= concentration of A at time=0 k= rate constant t= time
(first-order rate laws) the integrated rate equation describes …, another form includes:
exponential decay;
ln[A]= ln[A]0 - kt
(first-order rate laws) slope of the line=
y-intercept=
-k; [A]0
(first-order rate laws) a large value for k implies a
fast reaction
(first-order rate laws) half-life: the time needed for the concentration of a reactant to decrease to
1//2 its original value
(first-order rate laws) a short half-life indicates a … reaction
rapid
(first-order rate laws) for first-order, t1/2= …. and is …. of the concentration of the reactant
0.693/k; independent
(second-order rate laws) some reactions show rates that are proportional to the concentration of the reactant…
raised to the second power
(second-order rate laws) reaction rate=
-Δ[A]/ Δ t = k[A]2
(second-order rate laws) the units of a second order rate constant are
L/mol*s
(second-order rate laws) if the graph of …. vs. time is a straight line, the reaction obeys second-order kinetics
1/reactant concentration
(second-order rate laws) integrated form:
1/[A]= 1/[A]0 + kt
(second-order rate laws) t1/2=
1/k[A]
(second-order rate laws) slope=
k
collisions between molecules are necessary foe
reactions to occur
the rates of most reactions increase dramatically with
temperature
collision theory: basic assumption= molecules must … in order to react
collide
collision frequency: the number of … per second
moleclar collisions
collision frequency: Z=
Z0[A][B]
Z0 is a proportionality constant that depends on the … and … of A and B
speeds; sizes
experimental evidence shows that not every collision results
in a chemical reaction
activation energy: the minimum collision energy required ….
for a reaction to occur
activation energy: only collisions with enough energy to …. can result in the formation of products
rearrange bonds
activation energy: if the total energy of colliding species is too small, the molecules simply
bounce off each other
(activated complex→[X]*) intermediate species in chemical reactions are … than either the reactants or products of a chemical reaction
higher in energy
(activated complex→[X]*) the high-energy activated complex is very … and is also referrred to as the …
unstable; transition state
(activated complex→[X]*) the activation energy is the energy needed to form the … from the reactants
activated complex
the number of collisions with energies that exceed Ea grows exponentially with
temperature
the fraction of collisions with energy in excess of Ea can be derived from
fr= -e^Ea/RT
as t increases, fr … and approaches …
increases; 100%
predicted rate=
Z0[A][B]e^-Ea/RT
steric factor: not all collissions with energies that exceed Ea are productive;
the geometry of collisions must also be considered- not every collision occurs with the reactants in the correct …. to produce products
orientation
Rate= pZ0[A][B]e^-Ea/RT, where p is the…., Zo[A][B] is the …, and e^-Ea/RT is the …
steric factor; collision frequency; fraction exceeding Ea
k=
Ae^-Ea/RT
catalysts are substances that … the rate of reaction without … in the reaction
speed up; getting consumed
catalysts generally lower … and increase … at any given temperature. Some may increase the …
activation energy; reaction rate; steric factor
homogenous catalysts: present in the same …. as the reacting molecules
phase
heterogenous catalysts: present in a …. than the reacting molecules
different phase
enzymes: large molecules (usually …) that catalyze specific …
proteins; bioligical reactions
numerous gaseous reactions are catalyzed via absorption to the surface of a
solid metal catalyst
(heterogenous catalysis) step 1: … and … of the reactants
adsorption; activation
(heterogenous catalysis) step 2: migration of the adsorbed reactants to the
surface
(heterogenous catalysis) step 3: … of the adsorbed substances
reaction
(enzymes) enzymes interact with reactant molecules in a way that places them in the … to form the products;
can increase reaction rates up to ..
often end with an … suffix
correct geometry; 10^14; ase
reactions can be manipulated to decrease … and improve … if we know the appropriate reaction mechanism
side reactions; yields
a reaction mechanism is the sequence of … that lead from reactants to products
elementary steps
many reactions require multipe steps with …. that are produced, then consumed, during the course of the reaction
intermediates
sometimes more than one reaction mechanism is possible and further experimentation can help determine which mechanism is
most likely to occur
catalysts may allow for a different set of … to proceed as compared to those observed for an uncatalyzed reaction
elementary steps
for a reactio mechanism to be valid: the sum of the elementary steps must give the … for the reaction
the mechanism must agree with the …
overall balanced equation; experimentally determined rate law
unimolecualr reaction: involves a … molecule and a … rate law
single; first order
bimolecular reaction: involves … molecules and a …. rate law
two; second order
termolecular reaction: involves … molecules and a … rate law
three; third order
rate-limiting: … step in proposed reaction mechanism
slow
the rate-limiting step is the
rate-determing step
the experimental rate law must agree with the
rate-determining step
