Science and Technology news... :: View topic

Chris Gordon-Smith · Guest

The subject line summarises my question, but it may need a bit more
background.

BACKGROUND

I am developing a simulation model (SimSoup) that will include 'Molecules'
and 'Reactions' between them. The purpose of the model is to investigate
metabolic theories of the origin of life.

The Molecules and Reactions between them will be much simplified versions of
the real thing. For example, the Molecules will be made from six or less
types of Atom, and they will be two dimensional. I will model bonds and
bond energies.

I want to make Reaction Rates and Reaction constants dependent on bond
energies, and have so far used the Arrenhius equation:-
Rate Constant = A.exp(-Activation Energy / RT)

THE QUESTION

Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

As an example of what I mean for B, I am envisaging a situation where we
have two molecules that react very strongly (say sodium and water). If I
bring my two molecules close together at absolute zero, will they react?
The Arrenhius equation would suggest not (because RT would be zero), but is
it possible that the activation energy will be provided by electrostatic
forces rather than by thermal energy? In short, can one molecule 'tear a
piece off' the other even if thermal energy alone is not sufficient to form
an activated complex?

I would welcome any comments on this that will help make my model as
realistic as possible. My understanding of this area so far is at
http://www.simsoup.info/SimSoup/News_2006.html
(under the entry for 1 May 2006). The ideas described there are built into
the latest version of the model, but so far not in a way that represents
bonds explicitely.

Chris Gordon-Smith
www.simsoup.info

Fred Kasner · Guest

Chris Gordon-Smith wrote:
[quote]The subject line summarises my question, but it may need a bit more
background.

BACKGROUND

I am developing a simulation model (SimSoup) that will include 'Molecules'
and 'Reactions' between them. The purpose of the model is to investigate
metabolic theories of the origin of life.

The Molecules and Reactions between them will be much simplified versions of
the real thing. For example, the Molecules will be made from six or less
types of Atom, and they will be two dimensional. I will model bonds and
bond energies.

I want to make Reaction Rates and Reaction constants dependent on bond
energies, and have so far used the Arrenhius equation:-
Rate Constant = A.exp(-Activation Energy / RT)

THE QUESTION

Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

As an example of what I mean for B, I am envisaging a situation where we
have two molecules that react very strongly (say sodium and water). If I
bring my two molecules close together at absolute zero, will they react?
The Arrenhius equation would suggest not (because RT would be zero), but is
it possible that the activation energy will be provided by electrostatic
forces rather than by thermal energy? In short, can one molecule 'tear a
piece off' the other even if thermal energy alone is not sufficient to form
an activated complex?

I would welcome any comments on this that will help make my model as
realistic as possible. My understanding of this area so far is at
http://www.simsoup.info/SimSoup/News_2006.html
(under the entry for 1 May 2006). The ideas described there are built into
the latest version of the model, but so far not in a way that represents
bonds explicitely.

Chris Gordon-Smith
www.simsoup.info

[/quote]
Read the classic book by Linus Pauling "The Nature of the Chemical
Bond". If you are aiming so high as to produce a simplistic chemistry
then something a little more sophisticated arrow should be in your quiver.
FK

Bill Penrose · Guest

On Oct 5, 3:56 am, Chris Gordon-Smith <use.addr...@my.homepage> wrote:
[quote]
Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
[/quote]
That>s close. In order for two molecules to react, they have to
collide with enough kinetic energy to break bonds. The higher the
temperature, the more molecules will have the kinetic energy needed to
break those bonds. The distribution of kinetic energies in a
population of molecules is called the Boltzman distribution, a skewed
'bellshaped' curve, which depends on temperature.

The activation energy is the constant that makes the Arrhenius
equation fit, even though some molecules will react at lower energies,
and many more at higher energies.

In an exothermic reaction, you get some of that energy back, in that
the kinetic energy of the product molecules is greater than that of
the reacting molecules. Vice versa for endothermic.

As Fred implied, lots and lots of this kind of work has been done in
the past, and it>s just good science to look up prior work.

Dangerous Bill

Chris Gordon-Smith · Guest

Bill Penrose wrote:

[quote]On Oct 5, 3:56Â am, Chris Gordon-Smith <use.addr...@my.homepage> wrote:

Here is the question: What exactly is the activation energy for a
reaction in real chemistry? For example, is it:-

A) Â The sum of the bond energies of the bonds that must be broken to make
the reaction go

That>s close. In order for two molecules to react, they have to
collide with enough kinetic energy to break bonds. The higher the
temperature, the more molecules will have the kinetic energy needed to
break those bonds. The distribution of kinetic energies in a
population of molecules is called the Boltzman distribution, a skewed
'bellshaped' curve, which depends on temperature.

The activation energy is the constant that makes the Arrhenius
equation fit, even though some molecules will react at lower energies,
and many more at higher energies.

In an exothermic reaction, you get some of that energy back, in that
the kinetic energy of the product molecules is greater than that of
the reacting molecules. Vice versa for endothermic.

As Fred implied, lots and lots of this kind of work has been done in
the past, and it>s just good science to look up prior work.

Dangerous Bill
[/quote]
Thanks. The main point I wanted to confirm was that the bonds have to be
broken by kinetic energy (as opposed, for example, to electrostatic energy
in the electron shells). For practical purposes I think this means thermal
energy.

It was what I thought from looking at the Arrenhius equation, and from my
textbook (Atkins), but I wanted to be sure that I had not missed something.
(Perhaps I had a recollection of someone many years ago saying something
that suggested some other reaction mechanism.)

Chris Gordon-Smith
www.simsoup.info

Chris Gordon-Smith · Guest

Fred Kasner wrote:

[quote]Chris Gordon-Smith wrote:
The subject line summarises my question, but it may need a bit more
background.

BACKGROUND

I am developing a simulation model (SimSoup) that will include
'Molecules' and 'Reactions' between them. The purpose of the model is to
investigate metabolic theories of the origin of life.

The Molecules and Reactions between them will be much simplified versions
of the real thing. For example, the Molecules will be made from six or
less types of Atom, and they will be two dimensional. I will model bonds
and bond energies.

I want to make Reaction Rates and Reaction constants dependent on bond
energies, and have so far used the Arrenhius equation:-
Rate Constant = A.exp(-Activation Energy / RT)

THE QUESTION

Here is the question: What exactly is the activation energy for a
reaction in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

As an example of what I mean for B, I am envisaging a situation where we
have two molecules that react very strongly (say sodium and water). If I
bring my two molecules close together at absolute zero, will they react?
The Arrenhius equation would suggest not (because RT would be zero), but
is it possible that the activation energy will be provided by
electrostatic forces rather than by thermal energy? In short, can one
molecule 'tear a piece off' the other even if thermal energy alone is not
sufficient to form an activated complex?

I would welcome any comments on this that will help make my model as
realistic as possible. My understanding of this area so far is at
http://www.simsoup.info/SimSoup/News_2006.html
(under the entry for 1 May 2006). The ideas described there are built
into the latest version of the model, but so far not in a way that
represents bonds explicitely.

Chris Gordon-Smith
www.simsoup.info

Read the classic book by Linus Pauling "The Nature of the Chemical
Bond". If you are aiming so high as to produce a simplistic chemistry
then something a little more sophisticated arrow should be in your quiver.
FK
[/quote]
Thanks. I see that it is available at Amazon - although at a price!

Chris Gordon-Smith
www.simsoup.info

Bob · Guest

On Sun, 05 Oct 2008 11:56:29 +0100, Chris Gordon-Smith
<use.address@my.homepage> wrote:

[quote]The subject line summarises my question, but it may need a bit more
background.

BACKGROUND

I am developing a simulation model (SimSoup) that will include 'Molecules'
and 'Reactions' between them. The purpose of the model is to investigate
metabolic theories of the origin of life.

The Molecules and Reactions between them will be much simplified versions of
the real thing. For example, the Molecules will be made from six or less
types of Atom, and they will be two dimensional. I will model bonds and
bond energies.

I want to make Reaction Rates and Reaction constants dependent on bond
energies, and have so far used the Arrenhius equation:-
Rate Constant = A.exp(-Activation Energy / RT)

THE QUESTION

Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

[/quote]
As a practical matter, it is empirical. And it varies. Catalysts lower
the activation energy.

This is important to your project. Life is based on catalyzed
reactions -- with the catalysts controlling what reactions occur.

How important was catalysis in the origin of life? Beats me. Some
propose that life began, in some sense, as precursors came together on
mineral surfaces. Those surfaces can increase local concentrations,
and also provide catalysis, lowering the activation energy. There are
people working on the topic, and my vague recollection is that they
have models. Surely they have taken activation energy into account.
Are you up on all that line of work?

[quote]As an example of what I mean for B, I am envisaging a situation where we
have two molecules that react very strongly (say sodium and water). If I
bring my two molecules close together at absolute zero, will they react?
The Arrenhius equation would suggest not (because RT would be zero), but is
it possible that the activation energy will be provided by electrostatic
forces rather than by thermal energy? In short, can one molecule 'tear a
piece off' the other even if thermal energy alone is not sufficient to form
an activated complex?
[/quote]

Sure. That is what catalysts do.

I see a couple of replies already. Both are from regulars here, who
consistently post good stuff. Somehow, catalysis never came up. Not
sure why. Maybe raising it is good. Others can jump in again if they
have more on this.

bob

number6 · Guest

On Oct 5, 5:56 am, Chris Gordon-Smith <use.addr...@my.homepage> wrote:
[quote]The subject line summarises my question, but it may need a bit more
background.

[/quote]
Since you>re getting serious answers ... I>ll just add ... When you
wake up in the morning ... and it feels so good ... that kick that
gets you out of bed ... is the activation energy ...

Chris Gordon-Smith · Guest

Bob wrote:

[quote]On Sun, 05 Oct 2008 11:56:29 +0100, Chris Gordon-Smith
use.address@my.homepage> wrote:

Thanks for your comments. I>ve added some responses below
The subject line summarises my question, but it may need a bit more
background.

snip background stuff

THE QUESTION

Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

As a practical matter, it is empirical. And it varies. Catalysts lower
the activation energy.
[/quote]
My understanding is that this is a way of thinking about it if we are
considering the overall reaction, but that the picture is a little
different if we consider it at the level of elementary reactions.

For example, suppose we have an overall reaction:-

A + B --> C

and that it is catalysed by X. One might say that X lowers the activation
energy. But if we look at it at the level of elementary raections, we might
have:-

i) A + X --> I1
ii) I1 + B --> I2
iii) I2 --> C + X

Here, I1 and I2 are intermediates. There are three separate activation
energies, one for each elementary reaction.

The way I look at it is that X didn>t really lower the activation energy, it
just made possible a reaction pathway where the activation energy for each
of the elementary reactions is lower than that for the 'direct'
(non-catalytic) reaction.

[quote]
This is important to your project. Life is based on catalyzed
reactions -- with the catalysts controlling what reactions occur.
[/quote]
It certainly is important, so I will welcome any comment on my understanding
(above) of the basic mechanism for catalysis.

[quote]
How important was catalysis in the origin of life? Beats me. Some
propose that life began, in some sense, as precursors came together on
mineral surfaces. Those surfaces can increase local concentrations,
and also provide catalysis, lowering the activation energy. There are
people working on the topic, and my vague recollection is that they
have models. Surely they have taken activation energy into account.
Are you up on all that line of work?
[/quote]
I think you may be talking about Gunter Wachtershauser>s Iron-Sulphur world
theory. This envisages a 'pioneer organism' with a mineral substructure and
an organic superstructure. The organic compounds arise from redox reactions
involving very simple molecules such as CO, CO2, H2S, N2 and HCN. The
theory differs substantially from the more widely known RNA world theory in
that it does not require any organic molecules to get started, and also has
the great advantage that it does not need to solve the problem of getting
complex molecules like RNA to replicate accurately.

[quote]

As an example of what I mean for B, I am envisaging a situation where we
have two molecules that react very strongly (say sodium and water). If I
bring my two molecules close together at absolute zero, will they react?
The Arrenhius equation would suggest not (because RT would be zero), but
is it possible that the activation energy will be provided by
electrostatic forces rather than by thermal energy? In short, can one
molecule 'tear a piece off' the other even if thermal energy alone is not
sufficient to form an activated complex?

Sure. That is what catalysts do.
[/quote]
I would really like to get to the bottom of this, because its critical to my
project (as you have also indicated above).

If a catalyst can 'tear a piece off' another molecule, then it should be
able to do this just by coming into contact with the molecule, regardless
of the temperature. But this is not what the Arrenhius equation says.

Arrenhius' equation is k = A.exp(-Ea/RT),

and this seems to say that if the temperature is zero, the rate constant
will be zero and the reaction will not go. The equation seems to be saying
that it is purely thermal (kinetic) energy that must overcome the
activation energy.

Have I misunderstood something?

[quote]

I see a couple of replies already. Both are from regulars here, who
consistently post good stuff. Somehow, catalysis never came up. Not
sure why. Maybe raising it is good. Others can jump in again if they
have more on this.
[/quote]
Thanks. One other thing. I think that the points I have made above apply in
the case of reactions between molecules, but what about situations where
ions are involved. Does that change the picture?

Chris Gordon-Smith
www.simsoup.info

Craig · Guest

On Oct 8, 1:08 pm, Chris Gordon-Smith <use.addr...@my.homepage> wrote:
[quote]Bob wrote:
On Sun, 05 Oct 2008 11:56:29 +0100, Chris Gordon-Smith
use.addr...@my.homepage> wrote:

Thanks for your comments. I>ve added some responses below>>The subject line summarises my question, but it may need a bit more
background.

snip background stuff

THE QUESTION

Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

As a practical matter, it is empirical. And it varies. Catalysts lower
the activation energy.
[/quote]
Catalysis is less empirical than it once was, but theory was still a
long way from leading experiments, last I followed it.

At the simplest level, one could imagine breaking a bond to be the
activation energy for a reaction. (Gotta make room for the new bond
to come in.) However, in practice, this is almost always a
pessimistic overestimate of the activation energy. Most reactions do
not involve the energetically expensive cleavage of an otherwise
stable bond.

For example, consider the solution-phase reaction of, say, methyl
chloride with hydroxide to become methanol and free chloride. This
would NOT proceed by a mechanism like
CH3Cl --> [CH3]+ + Cl-
[CH3]+ + [OH]- --> CH3OH
It>s theoretically possible, but the first step is quite unfavorable
(i.e. very high activation energy). A more reasonable mechanism looks
like this:
[OH]- + CH3Cl --> [HO -CH3-Cl]-
[HO -CH3-Cl]- --> CH3OH + Cl-
In other words, the species associate with each other before any bonds
are broken. This requires much less energy, since bonds do not have
to be fully broken for the reaction to proceed. Most reaction
mechanisms that I am aware of involve similar kinds of "easing into"
the products from the reactants. The exceptions tend to be reactions
with particularly weak bonds (e.g. cleavage of iodine molecules to
iodine atoms at elevated temperatures). Dissociative mechanisms do
exist, for example
(CH3)3CBr --> [(CH3)3C]+ + Br-
[(CH3)3C]+ + OH- --> (CH3)3COH
However, these reactions are usually helped along by various ways of
stabilizing the intermediates (e.g. solvation). Again, the activation
energy is generally NOT equal to the gas phase bond cleavage energy.

[quote]The way I look at it is that X didn>t really lower the activation energy, it
just made possible a reaction pathway where the activation energy for each
of the elementary reactions is lower than that for the 'direct'
(non-catalytic) reaction.
[/quote]
I would say that a catalyst lowered the phenomenological activation
energy *because* it provided a new reaction pathway. :)

[quote]As an example of what I mean for B, I am envisaging a situation where we
have two molecules that react very strongly (say sodium and water). If I
bring my two molecules close together at absolute zero, will they react?
The Arrenhius equation would suggest not (because RT would be zero), but
is it possible that the activation energy will be provided by
electrostatic forces rather than by thermal energy? In short, can one
molecule 'tear a piece off' the other even if thermal energy alone is not
sufficient to form an activated complex?

Sure. That is what catalysts do.

I would really like to get to the bottom of this, because its critical to my
project (as you have also indicated above).

If a catalyst can 'tear a piece off' another molecule, then it should be
able to do this just by coming into contact with the molecule, regardless
of the temperature. But this is not what the Arrenhius equation says.

Arrenhius' equation is k = A.exp(-Ea/RT),

and this seems to say that if the temperature is zero, the rate constant
will be zero and the reaction will not go. The equation seems to be saying
that it is purely thermal (kinetic) energy that must overcome the
activation energy.

Have I misunderstood something?
[/quote]
Indeed, it is thermal energy that provides the "oomph" to climb the
activation energy barrier.

Most reactions have finite activation energies. Some reactions are
indeed so easy that they have essentially zero activation energy.
These usually involve very high energy species, like free atoms.
Consider this: anything so reactive as to have a very low barrier to
reaction would probably react very quickly. Most everyday/"kitchen"
chemistry is relatively boring, because all the truly reactive stuff
has already done its thing! Most biochemistry takes place in
relatively inert environments, where things like free atoms and sodium
metal have long since reacted with each other into less energetic
products. Situations like the upper atmosphere (constantly bombarded
with cosmic rays and UV light) and the early, pre-biotic Earth (hot;
lightning everywhere) allow more reactive species to exist.

Catalysts react "gently" with molecules to "coax" them into reacting
with each other. For catalysis to work, the catalyst has to gently
accept the reagents, allow them to react with each other, then let the
reaction products go. A great example of this is transition metal
heterogeneous catalysis. Metals far to the left (e.g. Ti) are so
electropositive that they are poor catalysts. They so desperately
want to bond with *something* that they never let anything go. For
example, titanium oxide is very difficult to decompose. At the other
end of the spectrum, gold is relatively inert. It is so unreactive
that it is one of the few elements that can be found pure in nature.
Gold is a poor catalyst because is just won>t "grab onto" much of
anything. For example, gold oxide doesn>t exist. Metals like
platinum and nickel are great catalysts for several reactions. They
adsorb reactants, but not too tightly. For example, platinum oxide
exists, but it can be decomposed with very modest effort. The
reactants of a reaction catalyzed by Pt form a "soup" of adsorbed
species, which can interact with each other. The adsorbed species are
held gently enough that the products can desorb, making room for more
reactants to repeat the cycle.

Where am I going with all this? Again, assuming the activation
energies represent bond energies is a gross overestimate of the
barrier to reaction. Or, put another way, it is very common to find
pathways to reaction that do not require complete cleavage of one bond
to occur before any other bonds can form. Just to pluck some random
numbers, most single bonds have bond energies around 300-400 kJ/mol.
It is easy to find reactions with activation energies on the order of
50-100 kJ/mol, for example.

[quote]Thanks. One other thing. I think that the points I have made above apply in
the case of reactions between molecules, but what about situations where
ions are involved. Does that change the picture?
[/quote]
There is nothing particularly magical about whether species are ions
in solution or neutral molecules.

Having written all of this, it is not clear to me how much of this is
relevant to your project. I>m not sure what your goals are. If your
goal is to faithfully replicate the chemistry of specific systems, you
have a LONG way to go. If you want to get, say, iron-sulfide
chemistry right, you need to specifically model iron and sulfur in
great detail. However, it seems like your goal is to produce a model
system that is "like" chemistry, simply to suggest that certain
phenomena are plausible. In that case, it might be enough to
qualitatively model that reactions have finite activation energies
that are less than typical bond energies.

- Craig

Bob · Guest

On Wed, 08 Oct 2008 21:08:44 +0100, Chris Gordon-Smith
<use.address@my.homepage> wrote:

[quote]THE QUESTION

Here is the question: What exactly is the activation energy for a reaction
in real chemistry? For example, is it:-

A) The sum of the bond energies of the bonds that must be broken to make
the reaction go
B) Something else.

As a practical matter, it is empirical. And it varies. Catalysts lower
the activation energy.

My understanding is that this is a way of thinking about it if we are
considering the overall reaction, but that the picture is a little
different if we consider it at the level of elementary reactions.

For example, suppose we have an overall reaction:-

A + B --> C

and that it is catalysed by X. One might say that X lowers the activation
energy. But if we look at it at the level of elementary raections, we might
have:-

i) A + X --> I1
ii) I1 + B --> I2
iii) I2 --> C + X

Here, I1 and I2 are intermediates. There are three separate activation
energies, one for each elementary reaction.

The way I look at it is that X didn>t really lower the activation energy, it
just made possible a reaction pathway where the activation energy for each
of the elementary reactions is lower than that for the 'direct'
(non-catalytic) reaction.

[/quote]
Ok.

However, we often do not know the elementary reactions, so are
"forced" to deal with it at the "overall" level. The "effect" is to
lower Ea.

If you were dealing with Pt (or Ni) catalyzing addition of H2 to an
alkene, the key elementary reaction is the adsorption of H2 onto the
metal surface, which serves to stretch the H-H to 2 H. That>s stated
very informally, but I think it is the idea.

If you can work at that level, ok, but I am skeptical that you can.

[quote]
This is important to your project. Life is based on catalyzed
reactions -- with the catalysts controlling what reactions occur.

It certainly is important, so I will welcome any comment on my understanding
(above) of the basic mechanism for catalysis.

[/quote]
One thing I was taught long ago is that catalysis is usually empirical
-- mechanism not understood. I>m sure we understand more now, but I
fear it is still often true.

[quote]
How important was catalysis in the origin of life? Beats me. Some
propose that life began, in some sense, as precursors came together on
mineral surfaces. Those surfaces can increase local concentrations,
and also provide catalysis, lowering the activation energy. There are
people working on the topic, and my vague recollection is that they
have models. Surely they have taken activation energy into account.
Are you up on all that line of work?

I think you may be talking about Gunter Wachtershauser>s Iron-Sulphur world
theory. This envisages a 'pioneer organism' with a mineral substructure and
an organic superstructure. The organic compounds arise from redox reactions
involving very simple molecules such as CO, CO2, H2S, N2 and HCN. The
theory differs substantially from the more widely known RNA world theory in
that it does not require any organic molecules to get started, and also has
the great advantage that it does not need to solve the problem of getting
complex molecules like RNA to replicate accurately.

[/quote]
Hey, you were bold enough to tackle his name!

Ok, but my reason for bringing it in... I recall that he has done
modeling. If so, he has thought thru some of these issues. It does not
matter how close your proposed chem is to his, but it could be good if
you understood his "math". Obviously, you are free to "improve" it.
(Apparently, he made some famously incorrect predictions.) But it
could be a useful starting point. Perhaps there is "debate" of his
models in the lit, or models by others.

I do not know anything about that literature. I may have casually read
some of his stuff, but my most recent foray into the topic was Bob
Hazen>s wonderful book, Genesis. It does not have technical detail,
but it does have references -- which I presume you could get anyway.

...

[quote]
If a catalyst can 'tear a piece off' another molecule, then it should be
able to do this just by coming into contact with the molecule, regardless
of the temperature. But this is not what the Arrenhius equation says.

Arrenhius' equation is k = A.exp(-Ea/RT),

and this seems to say that if the temperature is zero, the rate constant
will be zero and the reaction will not go. The equation seems to be saying
that it is purely thermal (kinetic) energy that must overcome the
activation energy.

Have I misunderstood something?
[/quote]
I think you have pushed the eqn beyond its limits. In the real world,
Ea depends on T. And all sorts of things happen (or fail to happen :-)
) at 0 K. See my example above about hydrogenation.

The Arrh eqn is basically empirical. It often holds -- closely enough
to be useful.

[quote]

I see a couple of replies already. Both are from regulars here, who
consistently post good stuff. Somehow, catalysis never came up. Not
sure why. Maybe raising it is good. Others can jump in again if they
have more on this.

Thanks. One other thing. I think that the points I have made above apply in
the case of reactions between molecules, but what about situations where
ions are involved. Does that change the picture?
[/quote]

Hm, not sure that any of the laws of thermo ask whether something is
an ion. Of course, the energy terms may be different. I>m inclined for
now to put this question on the back burner.

bob

Chris Gordon-Smith · Guest

Thanks for this. I>ll follow up on a few of the points.

[quote]
THE QUESTION

Here is the question: What exactly is the activation energy for a
reaction in real chemistry? For example, is it:-

A) Â The sum of the bond energies of the bonds that must be broken to
make the reaction go
B) Â Something else.

As a practical matter, it is empirical. And it varies. Catalysts lower
the activation energy.

Catalysis is less empirical than it once was, but theory was still a
long way from leading experiments, last I followed it.

At the simplest level, one could imagine breaking a bond to be the
activation energy for a reaction. (Gotta make room for the new bond
to come in.) However, in practice, this is almost always a
pessimistic overestimate of the activation energy. Most reactions do
not involve the energetically expensive cleavage of an otherwise
stable bond.

For example, consider the solution-phase reaction of, say, methyl
chloride with hydroxide to become methanol and free chloride. This
would NOT proceed by a mechanism like
CH3Cl --> [CH3]+ + Cl-
[CH3]+ + [OH]- --> CH3OH
It>s theoretically possible, but the first step is quite unfavorable
(i.e. very high activation energy). A more reasonable mechanism looks
like this:
[OH]- + CH3Cl --> [HO -CH3-Cl]-
[HO -CH3-Cl]- --> CH3OH + Cl-
[/quote]
Interesting example. Using my very sketchy knowledge of valence, I would
interpret this as a case in which the carbon in the methyl chloride has
valence 4, and so initially is 'happy' to have four bonds. Then along comes
the OH, and the carbon is now trying to bond to an additional atom
(presumably the O). It is a kind of 'overloading', and something has to
give. The weakest point is the Cl bond, so off goes the Cl atom.

[quote]In other words, the species associate with each other before any bonds
are broken. This requires much less energy, since bonds do not have
to be fully broken for the reaction to proceed. Most reaction
mechanisms that I am aware of involve similar kinds of "easing into"
the products from the reactants. The exceptions tend to be reactions
with particularly weak bonds (e.g. cleavage of iodine molecules to
iodine atoms at elevated temperatures). Dissociative mechanisms do
exist, for example
(CH3)3CBr --> [(CH3)3C]+ + Br-
[(CH3)3C]+ + OH- --> (CH3)3COH
However, these reactions are usually helped along by various ways of
stabilizing the intermediates (e.g. solvation). Again, the activation
energy is generally NOT equal to the gas phase bond cleavage energy.

The way I look at it is that X didn>t really lower the activation energy,
it just made possible a reaction pathway where the activation energy for
each of the elementary reactions is lower than that for the 'direct'
(non-catalytic) reaction.

I would say that a catalyst lowered the phenomenological activation
energy *because* it provided a new reaction pathway. :)
[/quote]
Perhaps its just a difference in perspective. My model works on elementary
reactions, and so I tend to think in thise terms.

The reason for modelling at the level of elementary reactions is that I am
modeling a 'soup' with very many molecular species and many pathways. It is
possible that the intermediate for one overall reaction (R) is also a
participant in another reaction. In this situation, I cannot work out the
rate for reaction R just by knowing the concentrations of its reactants and
products. I also have to know how the concentrations of intermediates are
affected by other reactions.
[quote]

As an example of what I mean for B, I am envisaging a situation where
we have two molecules that react very strongly (say sodium and water).
If I bring my two molecules close together at absolute zero, will they
react? The Arrenhius equation would suggest not (because RT would be
zero), but is it possible that the activation energy will be provided
by electrostatic forces rather than by thermal energy? In short, can
one molecule 'tear a piece off' the other even if thermal energy alone
is not sufficient to form an activated complex?

Sure. That is what catalysts do.

I would really like to get to the bottom of this, because its critical to
my project (as you have also indicated above).

If a catalyst can 'tear a piece off' another molecule, then it should be
able to do this just by coming into contact with the molecule, regardless
of the temperature. But this is not what the Arrenhius equation says.

Arrenhius' equation is Â k = A.exp(-Ea/RT),

and this seems to say that if the temperature is zero, the rate constant
will be zero and the reaction will not go. The equation seems to be
saying that it is purely thermal (kinetic) energy that must overcome the
activation energy.

Have I misunderstood something?

Indeed, it is thermal energy that provides the "oomph" to climb the
activation energy barrier.

Most reactions have finite activation energies. Some reactions are
indeed so easy that they have essentially zero activation energy.
These usually involve very high energy species, like free atoms.
Consider this: anything so reactive as to have a very low barrier to
reaction would probably react very quickly. Most everyday/"kitchen"
chemistry is relatively boring, because all the truly reactive stuff
has already done its thing! Most biochemistry takes place in
relatively inert environments, where things like free atoms and sodium
metal have long since reacted with each other into less energetic
products. Situations like the upper atmosphere (constantly bombarded
with cosmic rays and UV light) and the early, pre-biotic Earth (hot;
lightning everywhere) allow more reactive species to exist.

Catalysts react "gently" with molecules to "coax" them into reacting
with each other. For catalysis to work, the catalyst has to gently
accept the reagents, allow them to react with each other, then let the
reaction products go. A great example of this is transition metal
heterogeneous catalysis. Metals far to the left (e.g. Ti) are so
electropositive that they are poor catalysts. They so desperately
want to bond with *something* that they never let anything go. For
example, titanium oxide is very difficult to decompose. At the other
end of the spectrum, gold is relatively inert. It is so unreactive
that it is one of the few elements that can be found pure in nature.
Gold is a poor catalyst because is just won>t "grab onto" much of
anything. For example, gold oxide doesn>t exist. Metals like
platinum and nickel are great catalysts for several reactions. They
adsorb reactants, but not too tightly. For example, platinum oxide
exists, but it can be decomposed with very modest effort. The
reactants of a reaction catalyzed by Pt form a "soup" of adsorbed
species, which can interact with each other. The adsorbed species are
held gently enough that the products can desorb, making room for more
reactants to repeat the cycle.

Where am I going with all this? Again, assuming the activation
energies represent bond energies is a gross overestimate of the
barrier to reaction. Or, put another way, it is very common to find
pathways to reaction that do not require complete cleavage of one bond
to occur before any other bonds can form. Just to pluck some random
numbers, most single bonds have bond energies around 300-400 kJ/mol.
It is easy to find reactions with activation energies on the order of
50-100 kJ/mol, for example.

Thanks. One other thing. I think that the points I have made above apply
in the case of reactions between molecules, but what about situations
where ions are involved. Does that change the picture?

There is nothing particularly magical about whether species are ions
in solution or neutral molecules.

Having written all of this, it is not clear to me how much of this is
relevant to your project. I>m not sure what your goals are. If your
goal is to faithfully replicate the chemistry of specific systems, you
have a LONG way to go. If you want to get, say, iron-sulfide
chemistry right, you need to specifically model iron and sulfur in
great detail. However, it seems like your goal is to produce a model
system that is "like" chemistry, simply to suggest that certain
phenomena are plausible. In that case, it might be enough to
qualitatively model that reactions have finite activation energies
that are less than typical bond energies.
It is definitely relevant that it is not simply a question of the bond[/quote]
energies, and that molecules are flexy / bendy things that can 'ease into'
reactions without directly breaking bonds. I>m not sure to what extent I
can represent this, or to what extent it is necessary.

I am certainly not trying to faithfully replicate real chemistry. I am
really just trying to get something that is 'like' chemistry. The existing
model represents metabolic networks of reactions between molecules that
have mass and energy, but no structure. The networks can evolve, carrying
inherited information in their composition, rather than in genes. (It turns
out that the networks are attractors, which gives them the stability to
carry inherited information accurately). I would like to see whether giving
the molecules some structure makes things more interesting. It should make
the evolution more open ended, because the system will construct its own
molecules (currently I pre-define what types of molecules there are).
Perhaps I will also see 'interesting' molecules developing if I select for
more and more efficient networks.

I don>t know to what extent I>ll be able to represent the 'flexy / bendy'
nature of real molecules. I think I>ll have to start with rigid (two
dimensional) molecules. However, it is very useful to understand a little
better the nature of real molecules and reactions, and therefore the
implications of the simplifications I am making. Perhaps I>ll be able to
build in the idea that in some cases bonds can effectively be weakened by
the addition of a new group to a molecule (which is the way I read the
effect of the addition of the OH group to CH3Cl in your example; it weakens
the C-Cl bond.)

Chris Gordon-Smith
www.simsoup.info

[quote]
- Craig[/quote]

Chris Gordon-Smith · Guest

Bob wrote:

[quote]
For example, suppose we have an overall reaction:-

A + B --> C

and that it is catalysed by X. One might say that X lowers the activation
energy. But if we look at it at the level of elementary raections, we
might have:-

i) A + X --> I1
ii) I1 + B --> I2
iii) I2 --> C + X

Here, I1 and I2 are intermediates. There are three separate activation
energies, one for each elementary reaction.

The way I look at it is that X didn>t really lower the activation energy,
it just made possible a reaction pathway where the activation energy for
each of the elementary reactions is lower than that for the 'direct'
(non-catalytic) reaction.

Ok.

However, we often do not know the elementary reactions, so are
"forced" to deal with it at the "overall" level. The "effect" is to
lower Ea.

If you were dealing with Pt (or Ni) catalyzing addition of H2 to an
alkene, the key elementary reaction is the adsorption of H2 onto the
metal surface, which serves to stretch the H-H to 2 H. That>s stated
very informally, but I think it is the idea.

If you can work at that level, ok, but I am skeptical that you can.

I do work at the level of elementary reactions, but I think it will be[/quote]
difficult to represent this kind of 'stretching'. I think I>ll have to
start with rigid molecules (at least initially) and accept that its a
simplification.
[quote]

This is important to your project. Life is based on catalyzed
reactions -- with the catalysts controlling what reactions occur.

It certainly is important, so I will welcome any comment on my
understanding (above) of the basic mechanism for catalysis.

One thing I was taught long ago is that catalysis is usually empirical
-- mechanism not understood. I>m sure we understand more now, but I
fear it is still often true.

How important was catalysis in the origin of life? Beats me. Some
propose that life began, in some sense, as precursors came together on
mineral surfaces. Those surfaces can increase local concentrations,
and also provide catalysis, lowering the activation energy. There are
people working on the topic, and my vague recollection is that they
have models. Surely they have taken activation energy into account.
Are you up on all that line of work?

I think you may be talking about Gunter Wachtershauser>s Iron-Sulphur
world theory. This envisages a 'pioneer organism' with a mineral
substructure and an organic superstructure. The organic compounds arise
from redox reactions involving very simple molecules such as CO, CO2, H2S,
N2 and HCN. The theory differs substantially from the more widely known
RNA world theory in that it does not require any organic molecules to get
started, and also has the great advantage that it does not need to solve
the problem of getting complex molecules like RNA to replicate accurately.

Hey, you were bold enough to tackle his name!
But I stopped short of the little dots over the 'u' and the 'a>s!

Ok, but my reason for bringing it in... I recall that he has done
modeling. If so, he has thought thru some of these issues. It does not
matter how close your proposed chem is to his, but it could be good if
you understood his "math". Obviously, you are free to "improve" it.
(Apparently, he made some famously incorrect predictions.) But it
could be a useful starting point. Perhaps there is "debate" of his
models in the lit, or models by others.
Wachtershauser tends to focus on specific chemical pathways that could have[/quote]
existed in a 'pioneer organism'. I don>t think he puts a great deal of
emphasis on a mathematical approach. Others do take a very mathematical
approach. For example, Jain and Krishna take a network theory based
approach to show the behaviour of metabolic networks. Krisna>s PhD thesis
is called "Formation and Destruction of Autocatalytic Sets in an Evolving
Network Model"

[quote]
I do not know anything about that literature. I may have casually read
some of his stuff, but my most recent foray into the topic was Bob
Hazen>s wonderful book, Genesis. It does not have technical detail,
but it does have references -- which I presume you could get anyway.

..

If a catalyst can 'tear a piece off' another molecule, then it should be
able to do this just by coming into contact with the molecule, regardless
of the temperature. But this is not what the Arrenhius equation says.

Arrenhius' equation is k = A.exp(-Ea/RT),

and this seems to say that if the temperature is zero, the rate constant
will be zero and the reaction will not go. The equation seems to be saying
that it is purely thermal (kinetic) energy that must overcome the
activation energy.

Have I misunderstood something?

I think you have pushed the eqn beyond its limits. In the real world,
Ea depends on T. And all sorts of things happen (or fail to happen :-)
) at 0 K. See my example above about hydrogenation.

The Arrh eqn is basically empirical. It often holds -- closely enough
to be useful.
OK - Thanks. Sounds like it will be good enough for the kinds of scenarios I[/quote]
want to model (I don>t need to go to absolute zero!)
[quote]

I see a couple of replies already. Both are from regulars here, who
consistently post good stuff. Somehow, catalysis never came up. Not
sure why. Maybe raising it is good. Others can jump in again if they
have more on this.

Thanks. One other thing. I think that the points I have made above apply
in the case of reactions between molecules, but what about situations
where ions are involved. Does that change the picture?

Hm, not sure that any of the laws of thermo ask whether something is
an ion. Of course, the energy terms may be different. I>m inclined for
now to put this question on the back burner.
Good. Looks like one less thing to worry about at this stage.[/quote]

Chris Gordon-Smith
www.simsoup.info


www.GetXFactor.com Leading Technology, Science, Agriculture News and information	Part of the Identityscape.com network... getxfactor.com jmoodmusic.com smartbusinesschoices.com mintdepot.com lowfaresalways.com evangelicalview.com shoppingpodder.com soproudlywehail.com webnews.ws currenthumor.com

Science and Technology news... Forum Index -> Chemistry Forum
Page 1 of 1