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Help ! (about Nernst equation)
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Evgenij Barsukov
science forum Guru Wannabe


Joined: 09 May 2005
Posts: 137

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Help ! (about Nernst equation) Reply with quote

Your considerations about Ks would apply to something insoluble like
AgCl, but FeCl3 is highly soluble in water so you can deal with
conventional concentrations.

However, Nernst equation is strictly true only for diluted solutions.
For concentrated solutions corrections apply, making things more
complicated. Also FeCl3 can hydrolyse with building stuff like Fe(OH)2+
etc. Degree of hydrolysis will indeed depend on pH.
Nerst will do for rough estimate.

Regards,
Yevgen

n2mp wrote:
Quote:
Hello everyone,

To calculate the Nersnt potential of a FeCl3 solution, you just need to
apply the formula :

Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln([Fe3+]/ [Fe2+]).

However, as far as I remember, FeCl3 exists in a solid compound that
dissolves in water with a solubility constant :
Ks = [Fe3+]*[Cl-]^3. And, in the Nernst formula, that's the concentration
(in mol/L) of free Fe3+ in the bulk solution that has to be taken into
account.

If the FeCl3 is not completely dissolved in the bath, the formula becomes :
Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln(Ks/([Fe2+]*[Cl-]^3)).

The potential of a FeCl3 solution has its own value. However, on a pratical
point of view, you can only measure a potential difference between two
electrodes (the one on which you make the reaction, in this case an inert
material) and a reference electrode.
By convention, all Nernst potential are established and reported in tables
versus the NHE. But all reference electrodes have their potential vs the NHE
clearly established and reported on tables. Thus, the values are just
shifted by the potential differences between NHE potential and reference
electrode potential. In above formula, this only affects the E0(Fe3+/Fe2+)
value. In tables, you'll only find the value versus NHE but, since you'll
also find the potential shift between NHE and your reference electrode,
you'll be able toi deduced the suited E0 value in your operating conditions.

As you can see above, the reaction involves no protons. Thus, the
equilibrium potential is independant of pH. For other reactions, such as
water reduction or oxydation, protons are involved in the reaction and
obviously the pH impacts on Nernst potential value.

Some more details here :
http://www.cp.umist.ac.uk/lecturenotes/Echem/Electricityfile.html
More deeply in the subject :
http://electrochem.cwru.edu/ed/encycl/index.html

Best regards.

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n2mp
science forum beginner


Joined: 02 May 2005
Posts: 4

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Help ! (about Nernst equation) Reply with quote

Hello everyone,

To calculate the Nersnt potential of a FeCl3 solution, you just need to
apply the formula :

Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln([Fe3+]/ [Fe2+]).

However, as far as I remember, FeCl3 exists in a solid compound that
dissolves in water with a solubility constant :
Ks = [Fe3+]*[Cl-]^3. And, in the Nernst formula, that's the concentration
(in mol/L) of free Fe3+ in the bulk solution that has to be taken into
account.

If the FeCl3 is not completely dissolved in the bath, the formula becomes :
Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln(Ks/([Fe2+]*[Cl-]^3)).

The potential of a FeCl3 solution has its own value. However, on a pratical
point of view, you can only measure a potential difference between two
electrodes (the one on which you make the reaction, in this case an inert
material) and a reference electrode.
By convention, all Nernst potential are established and reported in tables
versus the NHE. But all reference electrodes have their potential vs the NHE
clearly established and reported on tables. Thus, the values are just
shifted by the potential differences between NHE potential and reference
electrode potential. In above formula, this only affects the E0(Fe3+/Fe2+)
value. In tables, you'll only find the value versus NHE but, since you'll
also find the potential shift between NHE and your reference electrode,
you'll be able toi deduced the suited E0 value in your operating conditions.

As you can see above, the reaction involves no protons. Thus, the
equilibrium potential is independant of pH. For other reactions, such as
water reduction or oxydation, protons are involved in the reaction and
obviously the pH impacts on Nernst potential value.

Some more details here :
http://www.cp.umist.ac.uk/lecturenotes/Echem/Electricityfile.html
More deeply in the subject :
http://electrochem.cwru.edu/ed/encycl/index.html

Best regards.


--
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WayneL1189
science forum beginner


Joined: 25 Mar 2005
Posts: 7

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis Reply with quote

Thanks

<nagy@anl.gov> wrote in message
news:1107794645.718604.41630@l41g2000cwc.googlegroups.com...
Quote:

Wayne,

No, 1.23 is srtictly connected to water decomposition, nothing to do
with dendrites.

Good luck: Zoltan.

WayneL wrote:
Hi Zoltan

I understand that but if the conditions were correct e.g electrode
that
would produce dendrites such as tin or copper, then does this voltage
have
any connection the starting of a dendrite spur?



Wayne

nagy@anl.gov> wrote in message
news:1107541263.266535.48080@z14g2000cwz.googlegroups.com...

Wayne,

The 1.23 has nothing to do with dendritic growth, only with
electrolysis of water to produce oxygen and hydrogen gas.
To get dendritic growth, you have to deposit a metal, and the
potential
of metal deposition will depend on what metal, what solution from
etc,etc...


Guud luck: Z.N.
WayneL wrote:
Hi

If you stayed just below the value of 1.23V is it possible to
experience
dendritic growth or does this have a min-max voltage?


Wayne

nagy@anl.gov> wrote in message
news:1107287327.008060.321020@f14g2000cwb.googlegroups.com...
A) Is there a minimum required voltage?

Yes, Theoretically, 1.23 volts. Practically, probably more like
2
or
more.

B) Can the current be stopped/started cyclicly?

Yes.

c) Does pressure hinder/enhance the process?

Yes, it requres more work (voltage) to produce a gas under
pressure
(it
would have required work to compress it)


D) Does cycles matter in process?

I do not understand the question. You cannot use ac. but you van
shut
off and retsart the process.

Good luck: Z.N.


Back to top
nagy@anl.gov
science forum beginner


Joined: 18 Jul 2005
Posts: 16

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis Reply with quote

Wayne,

No, 1.23 is srtictly connected to water decomposition, nothing to do
with dendrites.

Good luck: Zoltan.

WayneL wrote:
Quote:
Hi Zoltan

I understand that but if the conditions were correct e.g electrode
that
would produce dendrites such as tin or copper, then does this voltage
have
any connection the starting of a dendrite spur?



Wayne

nagy@anl.gov> wrote in message
news:1107541263.266535.48080@z14g2000cwz.googlegroups.com...

Wayne,

The 1.23 has nothing to do with dendritic growth, only with
electrolysis of water to produce oxygen and hydrogen gas.
To get dendritic growth, you have to deposit a metal, and the
potential
of metal deposition will depend on what metal, what solution from
etc,etc...


Guud luck: Z.N.
WayneL wrote:
Hi

If you stayed just below the value of 1.23V is it possible to
experience
dendritic growth or does this have a min-max voltage?


Wayne

nagy@anl.gov> wrote in message
news:1107287327.008060.321020@f14g2000cwb.googlegroups.com...
A) Is there a minimum required voltage?

Yes, Theoretically, 1.23 volts. Practically, probably more like
2
or
more.

B) Can the current be stopped/started cyclicly?

Yes.

c) Does pressure hinder/enhance the process?

Yes, it requres more work (voltage) to produce a gas under
pressure
(it
would have required work to compress it)


D) Does cycles matter in process?

I do not understand the question. You cannot use ac. but you van
shut
off and retsart the process.

Good luck: Z.N.

Back to top
WayneL1189
science forum beginner


Joined: 25 Mar 2005
Posts: 7

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis Reply with quote

Hi Zoltan

I understand that but if the conditions were correct e.g electrode that
would produce dendrites such as tin or copper, then does this voltage have
any connection the starting of a dendrite spur?



Wayne


<nagy@anl.gov> wrote in message
news:1107541263.266535.48080@z14g2000cwz.googlegroups.com...
Quote:

Wayne,

The 1.23 has nothing to do with dendritic growth, only with
electrolysis of water to produce oxygen and hydrogen gas.
To get dendritic growth, you have to deposit a metal, and the potential
of metal deposition will depend on what metal, what solution from
etc,etc...


Guud luck: Z.N.
WayneL wrote:
Hi

If you stayed just below the value of 1.23V is it possible to
experience
dendritic growth or does this have a min-max voltage?


Wayne

nagy@anl.gov> wrote in message
news:1107287327.008060.321020@f14g2000cwb.googlegroups.com...
A) Is there a minimum required voltage?

Yes, Theoretically, 1.23 volts. Practically, probably more like 2
or
more.

B) Can the current be stopped/started cyclicly?

Yes.

c) Does pressure hinder/enhance the process?

Yes, it requres more work (voltage) to produce a gas under pressure
(it
would have required work to compress it)


D) Does cycles matter in process?

I do not understand the question. You cannot use ac. but you van
shut
off and retsart the process.

Good luck: Z.N.

Back to top
WayneL1189
science forum beginner


Joined: 25 Mar 2005
Posts: 7

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis Reply with quote

Hi Zoltan

I understand that but if the conditions were correct e.g electrode that
would produce dendrites such as tin or copper, then does this voltage have
any connection the starting of a dendrite spur?



Wayne

<nagy@anl.gov> wrote in message
news:1107541263.266535.48080@z14g2000cwz.googlegroups.com...
Quote:

Wayne,

The 1.23 has nothing to do with dendritic growth, only with
electrolysis of water to produce oxygen and hydrogen gas.
To get dendritic growth, you have to deposit a metal, and the potential
of metal deposition will depend on what metal, what solution from
etc,etc...


Guud luck: Z.N.
WayneL wrote:
Hi

If you stayed just below the value of 1.23V is it possible to
experience
dendritic growth or does this have a min-max voltage?


Wayne

nagy@anl.gov> wrote in message
news:1107287327.008060.321020@f14g2000cwb.googlegroups.com...
A) Is there a minimum required voltage?

Yes, Theoretically, 1.23 volts. Practically, probably more like 2
or
more.

B) Can the current be stopped/started cyclicly?

Yes.

c) Does pressure hinder/enhance the process?

Yes, it requres more work (voltage) to produce a gas under pressure
(it
would have required work to compress it)


D) Does cycles matter in process?

I do not understand the question. You cannot use ac. but you van
shut
off and retsart the process.

Good luck: Z.N.

Back to top
nagy@anl.gov
science forum beginner


Joined: 18 Jul 2005
Posts: 16

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis Reply with quote

Wayne,

The 1.23 has nothing to do with dendritic growth, only with
electrolysis of water to produce oxygen and hydrogen gas.
To get dendritic growth, you have to deposit a metal, and the potential
of metal deposition will depend on what metal, what solution from
etc,etc...


Guud luck: Z.N.
WayneL wrote:
Quote:
Hi

If you stayed just below the value of 1.23V is it possible to
experience
dendritic growth or does this have a min-max voltage?


Wayne

nagy@anl.gov> wrote in message
news:1107287327.008060.321020@f14g2000cwb.googlegroups.com...
A) Is there a minimum required voltage?

Yes, Theoretically, 1.23 volts. Practically, probably more like 2
or
more.

B) Can the current be stopped/started cyclicly?

Yes.

c) Does pressure hinder/enhance the process?

Yes, it requres more work (voltage) to produce a gas under pressure
(it
would have required work to compress it)


D) Does cycles matter in process?

I do not understand the question. You cannot use ac. but you van
shut
off and retsart the process.

Good luck: Z.N.
Back to top
WayneL1189
science forum beginner


Joined: 25 Mar 2005
Posts: 7

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis Reply with quote

Hi

If you stayed just below the value of 1.23V is it possible to experience
dendritic growth or does this have a min-max voltage?


Wayne

<nagy@anl.gov> wrote in message
news:1107287327.008060.321020@f14g2000cwb.googlegroups.com...
Quote:
A) Is there a minimum required voltage?

Yes, Theoretically, 1.23 volts. Practically, probably more like 2 or
more.

B) Can the current be stopped/started cyclicly?

Yes.

c) Does pressure hinder/enhance the process?

Yes, it requres more work (voltage) to produce a gas under pressure (it
would have required work to compress it)


D) Does cycles matter in process?

I do not understand the question. You cannot use ac. but you van shut
off and retsart the process.

Good luck: Z.N.
Back to top
13131
science forum beginner


Joined: 25 Mar 2005
Posts: 6

PostPosted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Help ! (about Nernst equation) Reply with quote

In fact, by applying the Nernst equation the simple way I've got :
expr (0.77 + (0.0592 * 293.0 / 298.0 * (log(1.Cool+ 4.0)) - 0.2115) =
= 0.8255 V
at 21 C, and
[Fe2+] set to 10-6 M (just to divide by something > 0)
(unfortunately, I did not measure pH)
while my ORP electrode
(with a combined Ag/AgCl/(KCL 3.5M) internal reference) measured
0.6884 V
- quite a noticeable difference. Does
this make sense ? Will the required corrections account
for a less positive value - such as the measured one,
or are me and my ORP electrode doing something wrong ?
Thank
D.

Evgenij Barsukov wrote:

Quote:
Your considerations about Ks would apply to something insoluble like
AgCl, but FeCl3 is highly soluble in water so you can deal with
conventional concentrations.

However, Nernst equation is strictly true only for diluted solutions.
For concentrated solutions corrections apply, making things more
complicated. Also FeCl3 can hydrolyse with building stuff like Fe(OH)2+
etc. Degree of hydrolysis will indeed depend on pH.
Nerst will do for rough estimate.

Regards,
Yevgen

n2mp wrote:
Hello everyone,

To calculate the Nersnt potential of a FeCl3 solution, you just need to
apply the formula :

Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln([Fe3+]/ [Fe2+]).

However, as far as I remember, FeCl3 exists in a solid compound that
dissolves in water with a solubility constant :
Ks = [Fe3+]*[Cl-]^3. And, in the Nernst formula, that's the concentration
(in mol/L) of free Fe3+ in the bulk solution that has to be taken into
account.

If the FeCl3 is not completely dissolved in the bath, the formula becomes
: Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln(Ks/([Fe2+]*[Cl-]^3)).

The potential of a FeCl3 solution has its own value. However, on a
pratical point of view, you can only measure a potential difference
between two electrodes (the one on which you make the reaction, in this
case an inert material) and a reference electrode.
By convention, all Nernst potential are established and reported in
tables versus the NHE. But all reference electrodes have their potential
vs the NHE clearly established and reported on tables. Thus, the values
are just shifted by the potential differences between NHE potential and
reference electrode potential. In above formula, this only affects the
E0(Fe3+/Fe2+) value. In tables, you'll only find the value versus NHE
but, since you'll also find the potential shift between NHE and your
reference electrode, you'll be able toi deduced the suited E0 value in
your operating conditions.

As you can see above, the reaction involves no protons. Thus, the
equilibrium potential is independant of pH. For other reactions, such as
water reduction or oxydation, protons are involved in the reaction and
obviously the pH impacts on Nernst potential value.

Some more details here :
http://www.cp.umist.ac.uk/lecturenotes/Echem/Electricityfile.html
More deeply in the subject :
http://electrochem.cwru.edu/ed/encycl/index.html

Best regards.



--
__________________________________________
Back to top
Evgenij Barsukov
science forum Guru Wannabe


Joined: 09 May 2005
Posts: 137

PostPosted: Mon Mar 28, 2005 8:25 pm    Post subject: Re: Help ! (about Nernst equation) Reply with quote

As you have figured out because of your problems with [Fe2+],
Nernst equation is not very good to apply to situations very far
from 50/50 concentrations. Small changes in [Fe2+] have strong
effect on voltage. So, it does matter a lot what number you will set
there.
Because you can not set 0, you have to figure out what influences
this concentration. Mosty likely it is a reaction between
Fe(OH)3 that exists in equilibrium with oxigen, as
(1) Fe(OH)3 <--> Fe(OH)2 + 02 + H20.
Fe(OH)3 is itself built by hydrolysis of FeCl3 (2), so this
equilibrium will also play a role.

Basicaly, you are down to figuring out the exact concentration of
Fe2+, and this might well be an intractably complex theoretical task in
concentrated Fe3+ solution considering presence of above equilibriums.
This is even more difficult considering the fact that above mentioned
reactions involve solids and guesses, so their rate might be very slow.
Voltage might actually change after keeping your solution under open air
for a day or two.
It might be actualy easier to _measure_ Fe2+ concentration using
some complexing agent that makes colorful complex with Fe2+, and
measuring its concentration. Measuring pH would at least give you part
of the unknowns - degree of hydrolysis. But you still need to know
where equilibrium of (1) stands.

Another way to find out Fe2+ concentration is... did you guess?
Yes, voltage measurement, and "reverse engineering" the Nernst equation.
Just solve the reverse problem and find out what Fe2+ concentration will
give you the "right" potential.

Regards,
Evgenij

13131 wrote:

Quote:
In fact, by applying the Nernst equation the simple way I've got :
expr (0.77 + (0.0592 * 293.0 / 298.0 * (log(1.Cool+ 4.0)) - 0.2115) =
= 0.8255 V
at 21 C, and
[Fe2+] set to 10-6 M (just to divide by something > 0)
(unfortunately, I did not measure pH)
while my ORP electrode
(with a combined Ag/AgCl/(KCL 3.5M) internal reference) measured
0.6884 V
- quite a noticeable difference. Does
this make sense ? Will the required corrections account
for a less positive value - such as the measured one,
or are me and my ORP electrode doing something wrong ?
Thank
D.

Evgenij Barsukov wrote:


Your considerations about Ks would apply to something insoluble like
AgCl, but FeCl3 is highly soluble in water so you can deal with
conventional concentrations.

However, Nernst equation is strictly true only for diluted solutions.
For concentrated solutions corrections apply, making things more
complicated. Also FeCl3 can hydrolyse with building stuff like Fe(OH)2+
etc. Degree of hydrolysis will indeed depend on pH.
Nerst will do for rough estimate.

Regards,
Yevgen

n2mp wrote:

Hello everyone,

To calculate the Nersnt potential of a FeCl3 solution, you just need to
apply the formula :

Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln([Fe3+]/ [Fe2+]).

However, as far as I remember, FeCl3 exists in a solid compound that
dissolves in water with a solubility constant :
Ks = [Fe3+]*[Cl-]^3. And, in the Nernst formula, that's the concentration
(in mol/L) of free Fe3+ in the bulk solution that has to be taken into
account.

If the FeCl3 is not completely dissolved in the bath, the formula becomes
: Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln(Ks/([Fe2+]*[Cl-]^3)).

The potential of a FeCl3 solution has its own value. However, on a
pratical point of view, you can only measure a potential difference
between two electrodes (the one on which you make the reaction, in this
case an inert material) and a reference electrode.
By convention, all Nernst potential are established and reported in
tables versus the NHE. But all reference electrodes have their potential
vs the NHE clearly established and reported on tables. Thus, the values
are just shifted by the potential differences between NHE potential and
reference electrode potential. In above formula, this only affects the
E0(Fe3+/Fe2+) value. In tables, you'll only find the value versus NHE
but, since you'll also find the potential shift between NHE and your
reference electrode, you'll be able toi deduced the suited E0 value in
your operating conditions.

As you can see above, the reaction involves no protons. Thus, the
equilibrium potential is independant of pH. For other reactions, such as
water reduction or oxydation, protons are involved in the reaction and
obviously the pH impacts on Nernst potential value.

Some more details here :
http://www.cp.umist.ac.uk/lecturenotes/Echem/Electricityfile.html
More deeply in the subject :
http://electrochem.cwru.edu/ed/encycl/index.html

Best regards.



Back to top
13131
science forum beginner


Joined: 25 Mar 2005
Posts: 6

PostPosted: Tue Mar 29, 2005 10:55 am    Post subject: Re: Help ! (about Nernst equation) Reply with quote

I think I would be able to measure the actual concentration of Fe2+
by means of a photometric determination with 1,10-phenantroline.
This is quite easy.
D.

Evgenij Barsukov wrote:

Quote:
As you have figured out because of your problems with [Fe2+],
Nernst equation is not very good to apply to situations very far
from 50/50 concentrations. Small changes in [Fe2+] have strong
effect on voltage. So, it does matter a lot what number you will set
there.
Because you can not set 0, you have to figure out what influences
this concentration. Mosty likely it is a reaction between
Fe(OH)3 that exists in equilibrium with oxigen, as
(1) Fe(OH)3 <--> Fe(OH)2 + 02 + H20.
Fe(OH)3 is itself built by hydrolysis of FeCl3 (2), so this
equilibrium will also play a role.

Basicaly, you are down to figuring out the exact concentration of
Fe2+, and this might well be an intractably complex theoretical task in
concentrated Fe3+ solution considering presence of above equilibriums.
This is even more difficult considering the fact that above mentioned
reactions involve solids and guesses, so their rate might be very slow.
Voltage might actually change after keeping your solution under open air
for a day or two.
It might be actualy easier to _measure_ Fe2+ concentration using
some complexing agent that makes colorful complex with Fe2+, and
measuring its concentration. Measuring pH would at least give you part
of the unknowns - degree of hydrolysis. But you still need to know
where equilibrium of (1) stands.

Another way to find out Fe2+ concentration is... did you guess?
Yes, voltage measurement, and "reverse engineering" the Nernst equation.
Just solve the reverse problem and find out what Fe2+ concentration will
give you the "right" potential.

Regards,
Evgenij

13131 wrote:

In fact, by applying the Nernst equation the simple way I've got :
expr (0.77 + (0.0592 * 293.0 / 298.0 * (log(1.Cool+ 4.0)) - 0.2115) =
= 0.8255 V
at 21 C, and
[Fe2+] set to 10-6 M (just to divide by something > 0)
(unfortunately, I did not measure pH)
while my ORP electrode
(with a combined Ag/AgCl/(KCL 3.5M) internal reference) measured
0.6884 V
- quite a noticeable difference. Does
this make sense ? Will the required corrections account
for a less positive value - such as the measured one,
or are me and my ORP electrode doing something wrong ?
Thank
D.

Evgenij Barsukov wrote:


Your considerations about Ks would apply to something insoluble like
AgCl, but FeCl3 is highly soluble in water so you can deal with
conventional concentrations.

However, Nernst equation is strictly true only for diluted solutions.
For concentrated solutions corrections apply, making things more
complicated. Also FeCl3 can hydrolyse with building stuff like Fe(OH)2+
etc. Degree of hydrolysis will indeed depend on pH.
Nerst will do for rough estimate.

Regards,
Yevgen

n2mp wrote:

Hello everyone,

To calculate the Nersnt potential of a FeCl3 solution, you just need to
apply the formula :

Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln([Fe3+]/ [Fe2+]).

However, as far as I remember, FeCl3 exists in a solid compound that
dissolves in water with a solubility constant :
Ks = [Fe3+]*[Cl-]^3. And, in the Nernst formula, that's the
concentration (in mol/L) of free Fe3+ in the bulk solution that has to
be taken into account.

If the FeCl3 is not completely dissolved in the bath, the formula
becomes
: Eq(Fe3+/Fe2+) = E0 (Fe3+/Fe2+) + RT/F*ln(Ks/([Fe2+]*[Cl-]^3)).

The potential of a FeCl3 solution has its own value. However, on a
pratical point of view, you can only measure a potential difference
between two electrodes (the one on which you make the reaction, in this
case an inert material) and a reference electrode.
By convention, all Nernst potential are established and reported in
tables versus the NHE. But all reference electrodes have their potential
vs the NHE clearly established and reported on tables. Thus, the values
are just shifted by the potential differences between NHE potential and
reference electrode potential. In above formula, this only affects the
E0(Fe3+/Fe2+) value. In tables, you'll only find the value versus NHE
but, since you'll also find the potential shift between NHE and your
reference electrode, you'll be able toi deduced the suited E0 value in
your operating conditions.

As you can see above, the reaction involves no protons. Thus, the
equilibrium potential is independant of pH. For other reactions, such as
water reduction or oxydation, protons are involved in the reaction and
obviously the pH impacts on Nernst potential value.

Some more details here :
http://www.cp.umist.ac.uk/lecturenotes/Echem/Electricityfile.html
More deeply in the subject :
http://electrochem.cwru.edu/ed/encycl/index.html

Best regards.





--
__________________________________________
Back to top
glen herrmannsfeldt
science forum beginner


Joined: 30 Apr 2005
Posts: 36

PostPosted: Tue May 10, 2005 5:41 pm    Post subject: Re: Carbon rods - electrical connections methods Reply with quote

Rob wrote:

Quote:
Can anyone please advise me of methods used to make electrical connections
to carbon rods?

The connections are in a conductivity probe and are signal connections (ie
not high power. The current method involves simply twisting TCW around the
rods tightly then soldering it, this is then potted - not really ideal.


The preferred way for any conductivity measurement with questionable
contacts is with a four lead system, sometimes called a four point probe.

If you want the rods as part of the circuit, put two wires onto each
rod, supply current through one pair, measure the voltage across the
other. The contacts to the rods won't affect the measurement.

If you don't want the rods as part of the measurement you need four
rods, one wire each. Again, two supply the current (the outer two) and
two measure the voltage (the inner two). There are geometry corrections
to make the conductivity come out right.

-- glen
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Dieter Britz
science forum beginner


Joined: 04 May 2005
Posts: 45

PostPosted: Wed May 11, 2005 4:56 am    Post subject: Re: Carbon rods - electrical connections methods Reply with quote

glen herrmannsfeldt wrote:
Quote:
Rob wrote:

Can anyone please advise me of methods used to make electrical
connections
to carbon rods?


The connections are in a conductivity probe and are signal connections
(ie
not high power. The current method involves simply twisting TCW around
the
rods tightly then soldering it, this is then potted - not really ideal.



The preferred way for any conductivity measurement with questionable
contacts is with a four lead system, sometimes called a four point probe.

If you want the rods as part of the circuit, put two wires onto each
rod, supply current through one pair, measure the voltage across the
other. The contacts to the rods won't affect the measurement.

If you don't want the rods as part of the measurement you need four
rods, one wire each. Again, two supply the current (the outer two) and
two measure the voltage (the inner two). There are geometry corrections
to make the conductivity come out right.

This must be a pretty old question, and this posting does not
address the question itself, i.e. how best to contact the carbon.
How about gluing the contact wires onto the C with electrically
conducting epoxy? You can buy it, loaded with silver particles to
make it conducting. We use this and it works well.

--
Dieter Britz, Kemisk Institut, Aarhus Universitet, Danmark.
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WAYNEL1
science forum beginner


Joined: 04 May 2005
Posts: 25

PostPosted: Wed May 18, 2005 8:05 am    Post subject: Re: Req PIC/micro controller that I would be able to program/control with labview tia sal Reply with quote

Try the Pico Tech Ph Kit at around £100 ($200), this includes PC
connection + software. They sell the electrode separately.

http://www.picotech.com/ph-measuring.html


WayneL
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ketjow
science forum beginner


Joined: 07 Jun 2005
Posts: 6

PostPosted: Wed Jun 08, 2005 7:56 pm    Post subject: Re: easy way to obtaian C6H5-OH Reply with quote

Quote:
3rd step
the electrolysis of HO-C6H4-COONa(aq)--->> C6H5-ONa(aq) + CO2(g)

Could U explain why eletrolysis provides CO2?

And how U checked the gasseous CO2?

Are U sure CO2 doesn't comes from the excess of NaHCO3?


ketjow
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