Activation energy: Mg + HCl
James Hudgell
Mark Tarka
>
> Where can I find this value??
I'm not supposed to do this, but I've fallen
on hard times. For 10K English Pounds I will
provide you with the number you require. Cash
in advance, please.
Mark (You're not a chemist, are you...fess up :-)
Bill Nelson
: Where can I find this value??
Probably in the CRC Chemistry and Physics Handbook.
--
Bill Nelson (bi...@peak.org)
Ian Gay
Bill Nelson <bi...@spock.peak.org> wrote:
>James Hudgell <Ja...@hudgells.demon.co.uk> wrote:
>: Where can I find this value??
>
>Probably in the CRC Chemistry and Physics Handbook.
>
*** To reply by e-mail, remove _nospam from address ***
James Hudgell
>CRC Chemistry and Physics Handbook.
give me some clues....
--
JAMES HUDGELL
Ja...@Hudgells.demon.co.uk
James Hudgell
It can't be difficult for one of you helpful lot to look it up in one of
your big books. Go on - do your good deed for the year. Please.....
--
JAMES HUDGELL
Ja...@Hudgells.demon.co.uk
SNUMBER6
>Please help me - I'm only a poor little A-Level student.
>
>It can't be difficult for one of you helpful lot to look it up in one of
>your big books. Go on - do your good deed for the year. Please.....
>--
Activation energy is not usually tabularized anywhere ... It is generally
determined experimentally by measuring the dependence of rate constants on the
temperature of a reaction ... Or that's how I did it way back when ... The sum
of the energies of the products minus the energies of the reactants is the
overall energy produced ... The path to that reaction determines the activation
energy required to get the reaction to proceed ... If you knew the rate
determining step for say a chlorination was a reaction like Cl2 ===> 2 Cl
...then the rest of the reaction went to completion rapidly, then the
activation energy would be the energy of that reaction ... In your case, one
could only speculate as to the mechanism of the reaction with the rate
determing step being an undefined activated energy state of H+ or Mg ...
Are you sure you are looking for the Activation Energy for the reaction ??? And
also information is missing like is the HCl aqueous or gaseous ... Is the
reaction from a theoretical basis ??? or else it would be effected by the
surface area of the Mg ...
I believe many here would be glad to help you ... but can't for those reasons
...
I hope what I supplied is at least somewhat helpful ... maybe you can rephrase
the question for us ???
In the Village ....
I am not a number ... I am a free man !!!!
James Hudgell
I am investigating the "kinetics" of the reaction between aqueous
hydrochloric acid and magnesium ribbon. Obviously this produces
hydrogen, and measuring the time to produce 24cm^3 of gas is the main
way with which I shall be calculating "rates" of the equation.
To find how rate depends on concentration I used the same mass of
magnesium ribbon (~0.024g) each time and varied concentration of HCl(aq)
[0.5M --> 4.0M, in steps of .5M]. I found that rate = k.[HCl]^2.
I also decided to attempt to find the activation energy. To do this, I
fixed [HCl] (0.5M) and the mass of Mg (.024g). I varied temperature from
296K ==> 333K. I then calculated the activation energy from a graph of
ln(1/time) against 1/temp. The gradient of this x R (8.314) produced an
activation energy of 46.3 J.mol^-1.K^-1
Any ideas if this is reasonable??
--
JAMES HUDGELL - YEAR 12
Dr. Challoner's Grammar School - http://www.challoners.com/
Amersham, Buckinghamshire, UK.
SNUMBER6
>Any ideas if this is reasonable??
>--
How dare you ask us questions here !!!??
Based on this post you are well qualified to answer all questions that get
posted here ... Since my degrees are in Kinetics and Mechanisms, I can assure
you that your methodology to determine the answer is beyond reproach ... It
appears to be an excellent piece of work ...You know all about the Arrhenius
Plot to obtain your answer ...
I am not sure of your answer is correct though ... My memories are in
kcal/mole/K ... where a rule of thumb is about 12 kcal/mol/K is the range for a
typical reaction but most that I worked with were in the range 30-60 kcal/mol/K
... Unless I have a mental block on units ( which is entirely possible) your
result is about 12 cal/mole/K not Kcal ... I'd check your units on your
Arrhenius Plot ... that 1/temp usually has a 10^3 factor that sometimes get
missed ...and the other axis is ln(k) (which is proportional to 1/time) but
failing that ...go with your results you certainly seem to know what you are
doing ...
Feel free to email me if I can be of further assistance ...
BTW, two excellent books on kinetics theory and mechanisms are Eyring and
Eyring on kinetics ... and Ed King's book about chemical reactions ... They are
no doubt sitting in some dusty corner of the school library ...
BillyFish
>Date: Sat, Jul 15, 2000 9:12 AM
>Message-id: <ypPS1HA5...@hudgells.demon.co.uk>
>
>My apologies for being vague.
>
>I am investigating the "kinetics" of the reaction between aqueous
>hydrochloric acid and magnesium ribbon. Obviously this produces
>hydrogen, and measuring the time to produce 24cm^3 of gas is the main
>way with which I shall be calculating "rates" of the equation.
>
bubbles generated by the gaseous evolution effect the reaction? Does the fact
that the rate goes as k[HCl]^2 indicate that there is no effect from a
gas-metal interface?
Bill
SNUMBER6
>I do not have expert knowledge of this subject. My question would be: How do
>bubbles generated by the gaseous evolution effect the reaction? Does the fact
>that the rate goes as k[HCl]^2 indicate that there is no effect from a
>gas-metal interface?
Quite the contrary but not a gas-metal interface except as the reaction
proceeds to completion, the way he was measuring the reaction rate was to time
how long a certain volume of gas was generated ... Since both reactants were
present in excess, he was not evaluating the reaction through completion ...
In heterogenous actions like these, the rate dependence with the metal is
effected only by the surface area of the metal ... since he always used the
same weight of metal from the same presumed lot, this was a constant ... and
the k(observed) he found is = k'{Mg} ... the second order with respect to
hydrogen has mechanistic implications that suggest the presence of an
ion/metal/ion activated complex ...here the 2 electrons are rapidly transferred
from the Mg to the H+ putting the Mg++ into solution and hydrogen escaping
...So the gas/interface occurs after the rate determining step and is present
for only a very short time ...
BillyFish
>Date: Sat, Jul 15, 2000 5:42 PM
>Message-id: <20000715204201...@ng-fr1.aol.com>
Maybe I was not clear or did not understand the answer.
Doesn't the gas evolved reduce the effective surface area of the metal (Mg)
being eroded by the acid?
Bill
SNUMBER6
>Doesn't the gas evolved reduce the effective surface area of the metal (Mg)
>being eroded by the acid?
No ...actually as the Mg is eroded by the acid the surface area
increases(hopefully not enough to affect his rate) ... the gas formed is
present on the surface ( and blocking sites for attack as you state) for only a
very short period of time ... If this were significant, the reaction would NOT
be second order in HCl ...
James Hudgell
What causes a reaction to be first or second order? Is it to do with
rate determining steps?? How would this reaction progress? Is there some
kind of complex formed between H+ and Mg or the Cl- and the Mg or what??
Ideas??
Why is knowing the rate equation useful in "real life"?
--
JAMES HUDGELL - YEAR 12
Dr. Challoner's Grammar School - http://www.challoners.com/
Amersham, Buckinghamshire.
Mark Tarka
[good questions left for more knowledgeable folk]
> Why is knowing the rate equation useful in "real life"?
1. Knowing a rate equation for a specific reaction
can be the deciding factor when a chemical engineer
is considering a plant to produce a chemical product,
perhaps a fast reaction that works well in the lab on
a small scale, races to an unusable mess when tested
in a pilot plant (Bruce's P.A. reaction, for example :-)
2. Having lots of rate equations for lots of reactions
is known as a database from which inferences about an
unknown reaction might be tentatively drawn.
3. It's a great P. Chem lab project.
4. Rates of decomposition of biologically compatible
materials are factors in such things as breast implants,
artificial joints, subcutaneous time-released medication
devices.
5. Certainly leads to a firmer grasp on explosives, gun
powders, rocket fuels.
6. Organic chemists routinely control a synthesis step by
exploiting rates of reactions between certain chemicals
(kinetic vs. thermodynamic control) -- they use the concept
of a "rate equation" without ever knowing the acutal expression
and the values associated with it (have I got that correctly,
gentlemen?).
Anyone willing to provide more examples...?
Mark
SNUMBER6
>What causes a reaction to be first or second order?
God makes the rules ...we only investigate what they are :-) ...
The mechanism of the reaction determines what the reaction order is ... We use
the order and relationship to determine the mechanism ... By being second order
with respect to H+ implies that 2 protons must get together in the activated
complex formed in the rate determing step ...
>Is it to do with
>rate determining steps??
Definitely unless there are other things happening that are not known ...
>Is there some
>kind of complex formed between H+ and Mg or the Cl- and the Mg or what??
I would guess an activated complex {HMgH}*2+ is formed whereby two electrons
are transferred to the protons rapidly then the complex breaks apart into Mg2+
ions and H2 gas ...
Chloride is probably just along for the ride ... You can repeat with HClO4 as
the acid and see if there is a rate difference to prove this ...
Bufford L Hatchett
Would it be appropriate to snuh to this post?
JTTH...
B.L.H.
Dean Humphries
Snuh?
--
>>>> Get Snuhy!
>>>
>>> alt.tv.simpsons.snuh & alt.snuh.
>>
>> news:alt.snuh
>
> news:alt.tv.simpsons.snuh
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Cohort #15
M.I. #1138
Dr. Richard Silverł SPAMNEENERS CLUB Member
From: en_mog <en_...@hotmail.com>
Newsgroups: alt.tv.simpsons
Subject: Goodbye :-)
Date: Sat, 29 Apr 2000 08:05:45 -0500
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______
James Hudgell
>>From: bill...@aol.com (BillyFish)
>
>>I do not have expert knowledge of this subject. My question would be: How do
>>bubbles generated by the gaseous evolution effect the reaction? Does the fact
>>that the rate goes as k[HCl]^2 indicate that there is no effect from a
>>gas-metal interface?
>
>Quite the contrary but not a gas-metal interface except as the reaction
>proceeds to completion, the way he was measuring the reaction rate was to time
>how long a certain volume of gas was generated ... Since both reactants were
>present in excess, he was not evaluating the reaction through completion
with mine own eyes), the acid was in excess , but the magnesium was
being completely used up (0.001mol Mg ==> 24cm^3 of H2 = 0.001mol gas).
So there is an implication of changing surface area then...
>In heterogenous actions like these, the rate dependence with the metal is
>effected only by the surface area of the metal ... since he always used the
>same weight of metal from the same presumed lot, this was a constant
Initially yes, but obviously when it is all gone (==> MgCl2) then the
surface area is.... hmm
Mark Tarka
How about...
As the reaction proceeds, the mixing caused by the evolved
gas affects the rate of reaction...can you really characterize
the overall situation by using the surface area alone?
Classically...transport rates were the key...diffusion...etc.
Agitation by the product of a reaction (gas)...?
Please correct my useage of the vocabulary :-)
Mark (Let the FUNdamentals begin!)
BillyFish
>Date: Tue, Jul 18, 2000 1:45 PM
>Message-id: <BN5ecCAn...@hudgells.demon.co.uk>
>
>
>>>From: bill...@aol.com (BillyFish)
>>
>>>I do not have expert knowledge of this subject. My question would be:
>How do
>>>bubbles generated by the gaseous evolution effect the reaction? Does the
>fact
>>>that the rate goes as k[HCl]^2 indicate that there is no effect from a
>>>gas-metal interface?
>>
>>Quite the contrary but not a gas-metal interface except as the reaction
>>proceeds to completion, the way he was measuring the reaction rate was
>to time
>>how long a certain volume of gas was generated ... Since both reactants
>were
>>present in excess, he was not evaluating the reaction through completion
>
>As far as I thought (well in fact I know, because I saw it all disappear
>with mine own eyes), the acid was in excess , but the magnesium was
>being completely used up (0.001mol Mg ==> 24cm^3 of H2 = 0.001mol gas).
>So there is an implication of changing surface area then...
>
>>In heterogenous actions like these, the rate dependence with the metal
>is
>>effected only by the surface area of the metal ... since he always used
>the
>>same weight of metal from the same presumed lot, this was a constant
>Initially yes, but obviously when it is all gone (==> MgCl2) then the
>surface area is.... hmm
I thought that one point of limiting the gas to 24 ml was to stop the reaction
before there was much change in the surface area. As the Mg gets used up, the
surface area decreases, thereby reducing the rate,
Bill
James Hudgell
can look at it at:
http://www.hudgells.demon.co.uk/ChemInv/
I am sorry, but it is just produced form the M$ word / excel HTML
conversion wizards, so is not stylistically wonderful!!