Hello, in this video we're going to
start talking about chemical reactions. In this particular first video
of this module we're going to specifically talk
about what a reaction is, reaction rates, and
the reaction rate constant. So what is a chemical reaction? Well, it's a collisional process in which
the atomic composition of atoms and molecules change and
here is an example of such a reaction. CH3 is called the methyl radical, hydrogen atoms called
the hydrogen radical. Both of these are so-called
radical species and we'll get into that in a bit and
M could be any other molecule. So in this particular example, if the
reaction goes from the left to the right, it means that the methyl radical and
hydrogen collide, they form methane. But it just turns out that in this
kind of reaction if there was no way to absorb some of the extra
exothermic energy that that reaction involves then they just
immediately fall apart so that M could be any other molecule
that just absorbs some of that energy. And so
the products would be methane CH4 and the unchanged third body or the reaction could go
in the other direction. You could have methane and
it could be hit by some other molecule and it could knock a hydrogen off. In any event, the result of
the reaction is there's a shift in the distribution of the atoms,
but the atoms are conserved. So and this is reflected in
the coefficients in the equation. This is this equation has 1s in front
of each of the chemical symbols, but clearly if you start out with
one carbon on the left side, you have to end up with one
carbon on the right side. And that's reflected in the little
table that lists the carbon, hydrogen, and M,
whatever M may be, balance. And so the table lists the atom and
it lists the number in the reactants and the number in the products and
they have to be the same. There's 1 carbon, 4 hydrogens and 1 M, whatever M may be, and so they do balance. If you count up on the left
side the hydrogen, there's 3 in the CH3 and
then 1 that's free so that's 4. And on the right hand side there's
4 in the methane molecule. Now it's important to recognize that one sees two types of
chemical reactions written. The first is what's called
an elementary reaction. An elementary reaction is one that
actually occurs in a single step. So the example here is O atom
colliding with hydrogen molecule and knocking one of the hydrogens
off of the H2 and oxygen attaching in its place,
so you'll form + H. Global reactions
are reactions we often write. In undergraduate thermo,
if you've got to combustion, you will written an equation similar
to the one on the bottom of the slide. That's CH4 + 2O2 goes to CO2 + 2H2O. That's the global overall reaction of
the oxidation of methane with oxygen. And again,
if we count there's 1 carbon on each side, there's 4 hydrogens on each side and
there's 4 oxygen atoms on each side. Here are some additional
examples of elementary reactions. The first is a unimolecular
decomposition of acetylene. And that's where acetylene
falls apart into CH2 + C. In a more formal way of writing that
reaction there'd be an M on each side because of course acetylene isn't
going to fall apart all by itself. It's going to get whacked
by some other molecule and that will cause a carbon to detach. The next reaction is called
a bimolecular reaction because there's two reactants involved. In this example there's an O atom and
H2 and this is essentially this is the same reaction I showed
on the previous slide where O and H2 collide and
one of the Hs is traded out for an O. And then the third example is so
called trimolecular reaction where you have three reactants that come together,
CH + CO + M produces HCCO + M. So there's unimolecular reactions,
bimolecular reactions, and trimolecular reactions. You might notice that in
the trimolecular reaction the reaction that goes from right to left, M colliding with HCCO,
that is a unimolecular reaction. And that is often the case that
the two different directions, one direction of a unimolecular
reaction is unimolecular, the other direction would be trimolecular. And so those are the typical
kinds of reactions that you see. Now those are reactions,
what about how fast reactions occur? How do we express that? Well suppose we have aA
+bB goes to cC + dD. They have some coefficients in front of
them that reflects their composition and make sure that the number of atoms on
the left equals the number of atoms of the on the right of each type. Those lower case letters are called
stoichiometric coefficients. And it turns out that the overall
rate of reaction which is a concentration per unit time can be
written in the following fashion. Clearly if we write the overall
rate of reaction as the rate of going from the left to the right,
then A is decreased, B is decreased, C is increased,
D is increased. So we have R = -1 over a d[A]/dt or -1 over B, d[B]/dT or plus 1 over c d[C]/ dT or 1 over d d[D]/dt. Just note that the notation for
concentration is the molecular symbol
in square brackets. And that's amount of
material per unit volume and it could be number of
molecules per unit volume or moles of molecules per unit volume, but
it always on a number or molar basis. How about how fast the reaction occurs? Well, this is an expression that the slide
I just showed us gives an expression for the rate reaction but it doesn't say
much about how it actually goes. The reaction rate can be shown to usually depend on
the concentration of the reactants. Generally the reaction rate
is written as lowercase k which is called the reaction rate
constant times the concentration of A to some m power plus B to some n power. And if it's an elementary reaction,
that m and n are the stoichiometric
coefficients of A and B. And this is called and this produces what's called a rate
equation or rate expression. And you can write it symbolically
as the equation below. Again k is the reaction rate constant. So here's an example. Here's an elementary reaction, one that
I showed before, O + H2 goes to + H. There will be a reaction rate constant for each direction and I've labeled them
above and below the arrows here as kf forward and kr for reverse. Sometimes you'll see if you
have a number of reactions, say the first reaction might be k1 for
forward and k-1 for back. It's just a different notation. So then the rate expression for
say the destruction of O atom as a result of these reactions
d[O]/dt is going to be equal to minus k forward times the concentration of O
times the concentration of H2 plus k reverse times the concentration of
Times the concentration of hydrogen. Now often time especially in early days,
less so nowadays, where people didn't
understand very much about the details of the chemistry that took place in
these kinds of combustion reactions, they would write the global equation,
global reaction. So this is propane burned with oxygen. So we have propane + 5O2
produces 3CO2 + 4H2O. And then they would experimentally study
how fast this reaction occurred and they would come up with some
kind of rate expression for it. And this is an example so
that the rate of destruction of propane is equal to minus k times
the concentration of propane to the 0.1 times the concentration
of O2 to the 1.63. So what about the reaction rate constant? The reaction rate is determined by
the rate at which reactants collide times the probability that
a reaction will occur. And experiments show that many
rate constants are of the form shown where k is equal to some
constant a times the temperature to a power n times e to
the minus an energy over RT, the energy is called the activation
energy, R is the gas constant. T the temperature,
A is the pre exponential or frequency factor and often the value n is something other than than zero. So why does this make sense? Well, the activation energy is the energy
required for the reaction to occur, and we'll explore this in
more detail in a later video. But this little diagram shows
that if you have reactants that are of a certain energy and you want
to produce products you may have to provide an amount of energy
shown by this curve, which I've labeled the activation
energy for the reaction to occur. In other words, you have a molecule, it's
stuck together with inter-atomic forces, you've got to overcome that force in
order to pull the molecule apart and rearrange it with other atoms. Note that if the activation energy is
zero then that exponential term goes to zero and the rate is just some number and what that number is is
the collision frequency basically. Because it if the activation
energy is zero pretty much any collision can cause a reaction. In the world of chemical reaction
rates the way people study this is that they calculate the energy
of a particular molecule. One I've labeled A here, and then
they've calculated how much energy would be required in order for it to turn into
a different molecule B in this example. And then they produce a diagram
like this which is called a potential energy surface or
potential energy diagram. And so in this particular
example of A going to B, the activation energy is the difference
in energy of the A molecule and the energy of the highest
energy configuration that the molecule goes through in
converting to the B molecule. That's the activation energy. Now there is a simplification
in terms of calculating reaction rate constants that comes
about because of detailed balance. We talked about detailed balance
early earlier when we talk about chemical equilibrium and
the chemical equilibrium constant. It does turn out that if local
thermodynamic equilibrium applies, then detailed balance applies regardless
of whether there's chemical equilibrium. And what that means is that
the forward reaction rate constant and the reverse reaction rate
constants of an elementary reaction are related by a simple
function of temperature. And it turns out that that really
makes things a lot easier if you have a particular reaction and it's easier
to study in one direction than the other you study it in that direction and then you use detailed balance
to get the reverse rate. And also in computer codes, virtually all
the computer codes that are out there, you only have to provide the rate constant
information for one direction, and then they use detailed balanced
to get the other direction. So that's it for this video. Thanks for listening and have a great day.