NB. 201LM – lab prep on Wednesday 12  2
LECTURE 7 – Kinetics
Slide No. Notes
1how much product is formed
how much reactant disappears
2) actual ratio of reactant to product once reaction appears to stop
3)
2nothing to do with energy – kinetics
spontaneous because of thermodynamics
kinetics deals with how long it takes
3atmospheric chemistry – ozone depletion
biochemistry – medicines
archaeology – carbon dating
the more we understand of kinetics, more control over processes –
industrial reactions preferred to be fast – also slow down food
decomposition, aging
4more concentrated, more likely to collide earlier
more surface area, faster/more vigorous reaction
temperature
5red line = conc of A decreasing; green line – B conc increasing
disappearance of reactant or appearance of B
note negative sign
6 need to take into consideration coefficients
7Ans: B
Ans2: D
8note negative sign in average rate equation
9these are instantaneous rates
note that all three are same values
rate of reaction should be same for all reactants and products
do not matter which product or reactant examined, as long as
coefficients taken into consideration
LATER on, slope smaller, so reaction slows down
10 how rate changes with concentration
capitals are species of reaction; noncapitals are coefficients
double rate of N2O5, double rate of reaction – rate is proportional to
concentration
exponent = order of reaction

11 overall reaction order = sum of all exponents
Ans1: C
EX2.
double conc of A = increasing rate of reaction by four
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Halve conc of B = halve rate of reaction
Therefore Ans2: C
12 vary the initial concentration of one reactant at a time
set of experiments to measure rate and pick it apart by seeing how it
is affected by different reactant concentrations
why initial? Initially, there is no back reaction > no products, so
products cannot produce reactants and mess up concentrations
13 compare 1 and 2: ClO2 differs, so the difference between the
experiments’ initial rates is due to the ClO2
ClO2 – first order: both initial conc and initial rate increase by 4,
therefore onetoone therefore Ans1: B
Ans2: B
K = rate / [F2][ClO2] = 0.0012M/s divided by (0.010M)(0.010M) =
12 M^1s^1
Units depend on order of reaction
14 rate 3 divide by rate 2 – advantage: k and 0.02M^m cancels out
ditto for rate 1 divided rate 2
then plug in data from one experiment to find k
15 do not need to know integration of rate law
how product or reactant concentration changes with time
negative sign because writing concentration in terms of reactant
only for first order reaction
can use it calculate concentration at various times
16 exponential decay in first graph
know to be comfortable rearranging log equations
graphical order of determining rate of reaction
the slop is negative k
the y axis is In[A]
the yintercept is In[A]subscript0
17 half life independent of initial concentration
0.693 / k is constant
19 half life does depend on initial concentration
half life gets bigger as initial concentration decreases
20 independent of reactant concentration
yintercept is the initial concentration
21 ex of zeroorder reaction
only ammonia on surface of catalyst are reacting
on surface, rearrangement occurs so nitrogen and hydrogen forms
concentration independent – dependent on the surface
land on surface, nitrogen and hydrogen leave, more nitrogen land
22 order of reaction depends on mechanism, the individual steps of the
reaction
mostly 0, 1, and 2 order reactions in this course
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23 Ans: A which graph gives the straight line
B is not as tightly fitted as A
Ans2: given on slides
24 EX.
In[A]t – In[A]0 = kt
In[A]t – In(5.0x10^7M) = 1.45 y^1 (1yr)
 [A]t = 1.2x10^7 M
EX2.
In(3.0x10^7) – In(5.0x10^7) = 1.45y^1(t)
25  from molecular level of understand leads to understanding mechanism of
reaction
26 net reaction = sum of individual steps
N202 is both reactant and product so cancels out
Reaction intermediates never occur in reaction, just part of process
Termolecular is rare because statistically low to have three molecules
collide in correct orientation
27 rate depends directly on number of As and Bs because directly affects
number of collisions
28 EX.
2NO2Cl > 2NO2 + Cl2
step 1: unimolecular
step 2: bimolecular
rate = k[NO2Cl] for step 1
step 2: rate = k[NO2Cl][Cl]
k for both steps NOT THE SAME
all exponents one since no coefficients
Cl appears in rate law despite being intermediate
29 yes to first two questions, then F is the intermediate
SEE figure 12.33 on page 469
additional supporting evidence: presence of an intermediate
if predicted does not match with experiment, throw out the theory, not
the experiment
30 activation energy = mininum amount of energy for reaction to occur
see figure 12.16 on page 472
higher temperature – moving around faster – higher shift of
distribution – more successful collisions
32 sterics do matter
33 total energy is conserved
interplay between kinetic energy of reactants and products, and
potential energy
B in limbo, not completely belonging to either A or C
Transition state a very, very shortlived species
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