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Lecture

Notes taken during lecture


Department
Chemistry
Course Code
CHM135H1
Professor
Kris Quinlan

Page:
of 4
NB. 201LM – lab prep on Wednesday 12 - 2
LECTURE 7 – Kinetics
Slide No. Notes
1-how much product is formed
-how much reactant disappears
-2) actual ratio of reactant to product once reaction appears to stop
-3)
2-nothing to do with energy – kinetics
-spontaneous because of thermodynamics
-kinetics deals with how long it takes
3-atmospheric 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
4-more concentrated, more likely to collide earlier
-more surface area, faster/more vigorous reaction
-temperature
5-red 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
7-Ans: B
-Ans2: D
8-note negative sign in average rate equation
9-these 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; non-capitals 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 one-to-one 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 y-intercept 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
-y-intercept is the initial concentration
21 -ex of zero-order 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 short-lived species
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