Search   Memberlist   Usergroups
 Page 1 of 3 [38 Posts] View previous topic :: View next topic Goto page:  1, 2, 3 Next
Author Message
Evgenij Barsukov
science forum Guru Wannabe

Joined: 09 May 2005
Posts: 137

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Help ! (about Nernst equation)

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.
n2mp
science forum beginner

Joined: 02 May 2005
Posts: 4

 Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Help ! (about Nernst equation) 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. -- Enlevez ".nospam" de mon adresse e-mail pour me répondre. -------------------------------------------------------------------------------------------------------- Remove ".nospam" from my email address to reply me.
WayneL1189
science forum beginner

Joined: 25 Mar 2005
Posts: 7

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis

Thanks

<nagy@anl.gov> wrote in message
 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.
nagy@anl.gov
science forum beginner

Joined: 18 Jul 2005
Posts: 16

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis

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.
WayneL1189
science forum beginner

Joined: 25 Mar 2005
Posts: 7

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis

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
 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.
WayneL1189
science forum beginner

Joined: 25 Mar 2005
Posts: 7

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis

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
 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.
nagy@anl.gov
science forum beginner

Joined: 18 Jul 2005
Posts: 16

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis

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.
WayneL1189
science forum beginner

Joined: 25 Mar 2005
Posts: 7

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Electrolysis

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
 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.
13131
science forum beginner

Joined: 25 Mar 2005
Posts: 6

Posted: Fri Mar 25, 2005 7:41 am    Post subject: Re: Help ! (about Nernst equation)

In fact, by applying the Nernst equation the simple way I've got :
expr (0.77 + (0.0592 * 293.0 / 298.0 * (log(1.+ 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.

--
__________________________________________
Evgenij Barsukov
science forum Guru Wannabe

Joined: 09 May 2005
Posts: 137

Posted: Mon Mar 28, 2005 8:25 pm    Post subject: Re: Help ! (about Nernst equation)

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.+ 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.
13131
science forum beginner

Joined: 25 Mar 2005
Posts: 6

Posted: Tue Mar 29, 2005 10:55 am    Post subject: Re: Help ! (about Nernst equation)

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.+ 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.

--
__________________________________________
glen herrmannsfeldt
science forum beginner

Joined: 30 Apr 2005
Posts: 36

Posted: Tue May 10, 2005 5:41 pm    Post subject: Re: Carbon rods - electrical connections methods

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
Dieter Britz
science forum beginner

Joined: 04 May 2005
Posts: 45

Posted: Wed May 11, 2005 4:56 am    Post subject: Re: Carbon rods - electrical connections methods

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
make it conducting. We use this and it works well.

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

Joined: 04 May 2005
Posts: 25

 Posted: 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 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
ketjow
science forum beginner

Joined: 07 Jun 2005
Posts: 6

Posted: Wed Jun 08, 2005 7:56 pm    Post subject: Re: easy way to obtaian C6H5-OH

 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

 Display posts from previous: All Posts1 Day7 Days2 Weeks1 Month3 Months6 Months1 Year Oldest FirstNewest First
 Page 1 of 3 [38 Posts] Goto page:  1, 2, 3 Next View previous topic :: View next topic
 The time now is Mon Apr 22, 2019 10:08 am | All times are GMT
 Jump to: Select a forum-------------------Forum index|___Science and Technology    |___Math    |   |___Research    |   |___num-analysis    |   |___Symbolic    |   |___Combinatorics    |   |___Probability    |   |   |___Prediction    |   |       |   |___Undergraduate    |   |___Recreational    |       |___Physics    |   |___Research    |   |___New Theories    |   |___Acoustics    |   |___Electromagnetics    |   |___Strings    |   |___Particle    |   |___Fusion    |   |___Relativity    |       |___Chem    |   |___Analytical    |   |___Electrochem    |   |   |___Battery    |   |       |   |___Coatings    |       |___Engineering        |___Control        |___Mechanics        |___Chemical

 Topic Author Forum Replies Last Post Similar Topics Help me plaese with this equation.. Rjames2 Math 0 Fri Oct 13, 2006 3:23 pm Differential equation bamford Symbolic 0 Thu Aug 10, 2006 3:44 pm I need to know how it this equation was rearranged Alicia Math 3 Thu Jul 20, 2006 8:31 pm How many different ways of this this equation.... aliprinter Math 6 Mon Jul 10, 2006 10:48 am How to solve this PDE (laplace equation)? wandering.the.cosmos@gmai Research 0 Sun Jul 09, 2006 4:20 am