This course will cover the topics of a full year, two semester General Chemistry course. We will use a free on-line textbook, Concept Development Studies in Chemistry, available via Rice’s Connexions project. The fundamental concepts in the course will be introduced via the Concept Development Approach developed at Rice University. In this approach, we will develop the concepts you need to know from experimental observations and scientific reasoning rather than simply telling you the concepts and then asking you to simply memorize or apply them. So why use this approach? One reason is that most of us are inductive learners, meaning that we like to make specific observations and then generalize from there. Many of the most significant concepts in Chemistry are counter-intuitive. When we see where those concepts come from, we can more readily accept them, explain them, and apply them. A second reason is that scientific reasoning in general and Chemistry reasoning in particular are inductive processes. This Concept Development approach illustrates those reasoning processes. A third reason is that this is simply more interesting! The structure and reactions of matter are fascinating puzzles to be solved by observation and reasoning. It is more fun intellectually when we can solve those puzzles together, rather than simply have the answers to the riddles revealed at the outset. Recommended Background: The class can be taken by someone with no prior experience in chemistry. However, some prior familiarity with the basics of chemistry is desirable as we will cover some elements only briefly. For example, a prior high school chemistry class would be helpful. Suggested Readings: Readings will be assigned from the on-line textbook “Concept Development Studies in Chemistry”, available via Rice’s Connexions project. In addition, we will suggest readings from any of the standard textbooks in General Chemistry. A particularly good free on-line resource is Dickerson, Gray, and Haight, "Chemical Principles, 3rd Edition". Links to these two texts will be available in the Introduction module.
CSD 21 Reaction Kinetics I
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来自 Rice University 的课程
普通化学：基本概念与应用
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从本节课中
Chemical Kinetics
Chemical reactions occur at very different rates, some occurring so slowly that we only notice them with great passing of time and some occurring explosively rapidly. In this module, we develop measurements of the rates of reaction, determining the factors which can make a reaction proceed more rapidly or more slowly. These observations are summarized in equations called Rate Laws, where each reaction has its own empirical rate law. By using kinetic molecular theory, we develop a model to understand how and why each factor in the rate law is important in determining the rate of a chemical reaction.
与讲师见面
John Steven HutchinsonDean of Undergraduates and Professor of Chemistry
In the previous concept development studies, we began our understanding
of chemical reaction rates just from an empirical point of view.
Could we write down equations which related the rate
of the chemical reaction to, for example, the concentrations
of the reactants and to figure out, how then,
the concentration of the reactants would vary with time?
And the current concept development study over the course of the next
couple of lectures, we want to develop a model that helps us understand
why the reaction rate depends upon the concentrations in the way that it does.
Let's go back and take a look at that data again
on the experimentally observed rate laws for a whole collection of reactions.
You'll remember these data from before, where we
discovered, in fact, that there could be a
relationship between the stoichiometric coefficient of the materials
reacting, say, two nitric oxides reacting with a single
oxygen, and the exponents or the orders in the rate law for that reaction.
So for example, in this same reaction.
The exponent two on the nitric oxide matches the Stoichiometric
coefficient, two on the nitric oxide in the balanced equation.
And, an exponent on the oxygen one,
matches the Stoichiometric coefficient one on the oxygen.
But that doesn't always seem to be the case.
For example.
Two
ICLs reacting with one hydrogen, we'll notice that there isn't an alignment
between the stoichiometric coefficients on the ICL and the exponent on the ICL.
in a similar way, and a particularly complicated one, the reaction
of hydrogen and bromine even produces And a half integer, exponent.
But on some occasions, again, the last one here between ozone and chlorine,
there does seem to be an alignment between the stoichiometric coefficients
and the balanced equation, and the exponents in the rate law.
What we want to do is use this experimental data, to develop
a model that helps us understand, where do those variations come from.
And our first step in this process is to try
to begin to understand, how does a chemical reaction take place.
At the very lowest level, it seems pretty clear,
that if two molecules are to react with each other they have to collide.
They need to be in close proximity to one another for the atoms to be exchanged.
From one molecule to another. That seems straightforward.
So let's think about the probabilities of that event occurring.
We actually have looked at this before, for the case of an ideal
gas of say, molecules, A, flying about in random motion, in random directions.
We've studied this before.
And, we're asking the question, as these molecules
randomly move about, what is the probability that two
of them would essentially wind up in the
same time at a such that a collision occurs?
If they are confined to some space here, it's pretty clear that they are going to
run into each more frequently when there are
more of them, confine to a smaller space.
This is our own experience to sort of walking around in crowded venue
like a sports arena or walking around in uncrowded arenas like maybe a park.
The higher density of the people in the area,
the more likely you are to run into each other.
So the probability of a collision.
Depends upon how many particles are there in what volume.
The more particles there are per volume,
the greater will be the collision probability.
And in fact, since that is the
concentration, then the collision
probability depends upon the concentration.
Furthermore, It depends upon it in a way
which is slightly more complicated than just proportional.
Back when we studied the ideal gas law and the kinetic molecular theory, we came up
with an equation which actually described the rate
of collisions of A molecules with each other.
And notice the the things that it, that show up here.
There is the size
of the molecule, the d squared Tells us d is the diameter of the molecule.
This term here is the, essentially, average speed of the molecules.
It makes sense that the faster the molecules are
going, the more frequently they'll run into each other.
But mostly, and very importantly, notice that the dependence
upon the concentration depends upon the A particle's square.
That's because in order to have a collision say at
this particular point, I need to have two events occur simultaneously.
I need to have the probability of the A
molecule being here and a second A molecule being here.
So the concentration, I'm sorry, the the, collision probability goes up not
as the concentration of A, but as the concentration of A squared.
What if we want to have a reaction say
of A plus B running into each other. Now instead, I've got a container in which
I have mixed together B molecules and A molecules in some random fashion.
And the probability that an A molecule and
a B molecule in travelling intercept at some
particular point and have a collision, then depends
both upon the concentration of the A molecules and
the concentration of the B molecules.
The higher the density of B molecules, the more frequently they'll run into As.
But it also depends upon how many As there
are, the higher the concentration of the A molecules.
Now, we've already seen a little while ago that if we
had a reaction in which we were reacting, say, 2A molecules with
each other, the rate at which the reaction could occur could
not be any faster than the rate at which the A molecules
run into each other.
And we've already seen on the previous slide that that rate goes like A squared.
Notice an interesting trend here, that is, when the stoichiometric coefficient
on A and B is just 1, each of A and B are raised to the first power.
when the stoic humetric coefficient on A is 2, A is raised to the second power.
That actually suggests overall
that there's some kind of a relationship here, that the stoichiometry is related to
the order of the reaction rates, and
in particular the exponent of the reaction rate.
The rate law appears to be related to the stoichiometry of the collisions.
Sorry.
Is related to the stoichiometry of the collisions.
That's what comes
out of this collision theory model that we're developing here.
But now we go back and look at that data again.
Here are stoichiometry equations and here are rate laws.
And we've already observed earlier in this lecture that
they don't seem to be cleanly related to one another.
What that tells us is there must be something more to the story.
There must be something else going on
other than just molecules running into each other.
Apparently, other complexities must come in to play.
Let's consider then.
That what we've observed from our collision model is that if the
reaction occurs in a single step, say, A molecules must col, collide
with B molecules, or A molecules must collide with A molecules, then
the order of reaction is determined by the stoichiometry of the reaction.
But, since in many
cases, the order of the reaction does not
seem related to the stoichiometry of the collision,
we can conclude that the reaction then will
not occur, does not occur in a single collision.
If there's a single collision, the order in the stoichiometry are the same.
If the order in the stoichiometry are not the same,
then there isn't a single collision that produces the chemical reaction.
Let's do a specific example to make it clear what we are talking about here.
A specific example is going to be the reaction of nitrogen
dioxide with carbon monoxide to form nitric oxide and carbon dioxide.
The question is, does this reaction occur in a single collision?
Do we have, for example her, nitric
oxide molecules running into carbon dioxide molecules
in a single step?
It would seem like that would make sense, that
in order to exchange the oxygen from one molecule to
the other, All we would need to do is have
the nitrogen dioxide and the carbon monoxide get close together.
Let's test that model by looking at the rate law.
The experimentally observe rate law for this reaction does not depend upon
the concentration of the carbon monoxide molecules, and it depends doubly,
quadratically on the nitrogen dioxide molecule concentration.
So immediately you can conclude from the model we've
just developed, this reaction cannot occur in a single collision,
because if it did, the rate would depend upon
the concentration of the nitrogen dioxide to the first power.
And the concentration of the carbon monoxide to the
first power, and that's not what we observed experimentally.
Something else is going on,
so we can conclude that the reaction occurs in more than one step.
That leads us to the concept of a reaction mechanism.
A reaction mechanism is a series of
reactions, which lead to the overall reaction.
Each reaction in the series is a single collision, two molecules
running into each other and reacting
to produce some, perhaps, intermediate product.
And a series of these reactions were
collisions occurring in sequence, produces the overall reaction.
Let's give, let's see, an example of what this might look like.
So, since the rate law depends on the
nitrogen dioxide concentration squared, and in a single
collision, the rate will be given by the
concentrations of the products, or of the reactants together.
It must be true, in order to have this rate law, somewhere along the
line, the mechanism must include a reaction
step in which nitrogen dioxide molecules are running
into nitrogen dioxide molecules.
The rate of a process in which nitrogen dioxide molecules run into each other
would be nitrogen dioxide concentration squared times a proportionality constant.
And therefore we expect to have a mechanism in which that reaction occurs.
Here is a possible mechanism.
Imagine, what happens is first, two nitrogen
Dioxide molecules run into each other, producing two new materials here.
Notice what we wind up with is a NO3.
And in NO.
NO is in fact one of the products of the overall reaction.
NO3 is not.
So before the reaction is over, that NO3 is going to have to go away somewhere.
But we also have not yet used our carbon monoxide reactant, so perhaps
what happens is that the nitren nitrogen, the NO3.
Reacts with the carbon monoxide to produce finally
products of the reaction in O2 and CO2.
Notice that NO2 is produced, recreating one of
the reactants that could go back and can
participate in for the reaction, and the two
steps together, produce the two products of this reaction.
Nitric oxide and carbon dioxide.
Now, how would this mechanism be related to the overall rate law?
We know that the rate law has nitrogen dioxide
squared as its concentration, or, as its rate law.
And as a consequence, apparently that first step in
the process is more important then the second step.
As a consequence, we create a new model now in
which we say two steps take place here, but one
of those steps is much slower than the other step.
That's an interesting argument because it says I
have to accomplish two things before the reaction occur.
One of those things is very slow.
In fact I can go no faster than the slowest step in my process.
The slowest step in the process determines
my overall reaction rate, and as a consequence, we actually
refer to that slow step as the rate determining step.
That the rate of the reaction that must take place in the first step
of our overall reaction, that first step must determine the overall reaction rate.
And what is the speed of that first reaction?
Well, according to what we have done, it must be the
rate at which one nitrogen dioxide
collides with another dioxide, nitrogen dioxide.
That's exactly what we have done back over here
when we talked about A molecules running into one another.
Instead what we need, is nitrogen dioxide molecules running into each other.
And the rate in which nitrogen dioxide molecules run into each other, is going to
depend upon how many nitrogen dioxide molecules
there are per volume in this particular container.
That means it depends upon the concentration.
Of the nitrogen dioxide squared because the collision
rate of nitrogen dioxides with each other depends
upon the nitrogen dioxide squared.
Therefore, the rate of step one is k times NO2 squared.
And if step one is the slow step, the rate-determining step in the, in this
whole reaction, the overall rate of the reaction
is the rate of the slow moving step.
The other step occurs so rapidly that once the first step has
occurred, the second step immediately clicks in.
This, then, is an example of how a multi-step process can give rise
to an overall reaction and the slow step in that multi-step process determines what
the rate law is by determining what the stoichiometric coefficients in the rate
determining step are and thereby determining what
the rate is of the collisions in
that rate determining step.
We could develop mechanisms for a variety of different reactions, and each of
those mechanisms will help account for the
overall reaction rate law for that reaction.
There are other things that have to happen
besides collisions of molecules for reactions to take place.
We're going to consider those other characteristics in the next lecture.
