Sciencemadness Discussion Board - Quantitative test for chloride ion...

Quantitative test for chloride ion...

Swede - 16-9-2012 at 14:20

... that does NOT use silver nitrate. Is there such a beast?

I have to measure far too many aqueous solutions of chlorides, and I'd love to learn of a method that does not use expensive silver nitrate. I've been unable to find much using Google and the like.

When there is just one cation like Ca+ or Na+, sometimes that can be isolated in a precipitate and weighed, but when you have a mixture of differing chlorides, it can be a bit messy.

TIA and cheers

Vargouille - 16-9-2012 at 14:40

Just look up which other chlorides are very insoluble. To memory, that's PbCl2, CuCl, and Hg2Cl2. Of course, it's the Hg2Cl2 that's the best (Ksp of Hg2Cl2 is 1.3×10^-18), but I would guess that Hg2(NO3)2 is quite a bit more expensive than AgNO3. Cuprous salts are prone to oxidation, and from what I read, CuNO3 is not easily prepared. PbCl2 is still a bit soluble in water (Ksp=1.7×10^-5).

I looked up a bit more, and thallium(I) halides are also fairly insoluble, but they are also quite toxic, and TlCl is more soluble than PbCl2 (Ksp=1.86×10^-4).

So, for quantitative purposes, AgNO3 is probably your best bet.

EDIT: Color me impressed. Sigma-Aldrich sells >97% Hg2Cl2-2H2O for $73/50g and >99.0% AgNO3 for $97/25g. So, you might be able to use the Hg2(NO3)2 for your quantitative purposes, but the toxicity is still an issue. Purity may be an issue as well, but that's nothing a recrystallization couldn't fix, eh?

[Edited on 16-9-2012 by Vargouille]

[Edited on 16-9-2012 by Vargouille]

Lambda-Eyde - 16-9-2012 at 15:21

Quote: Originally posted by Vargouille

Just look up which other chlorides are very insoluble. To memory, that's PbCl2, CuCl, and Hg2Cl2. Of course, it's the Hg2Cl2 that's the best (Ksp of Hg2Cl2 is 1.3×10^-18), but I would guess that Hg2(NO3)2 is quite a bit more expensive than AgNO3. Cuprous salts are prone to oxidation, and from what I read, CuNO3 is not easily prepared. PbCl2 is still a bit soluble in water (Ksp=1.7×10^-5).

Translated to more familiar terms, a Ksp of 1.7e-5 for PbCl2 equals a solubility of 9.9 g/L at 20 °C, according to Wikipedia. Which is useless if you want to use it for quantitative determination of chloride. In comparison, again from Wiki:

Quote:

The solubility product, K_sp, for AgCl is 1.77 × 10^-10 M², which indicates that one liter of water will dissolve 1.3 × 10^-5 moles (1.9 mg) of AgCl at room temperature.

Quote: Originally posted by Vargouille

I looked up a bit more, and thallium(I) halides are also fairly insoluble, but they are also quite toxic, and TlCl is more soluble than PbCl2 (Ksp=1.86×10^-4).

No, it isn't; keep in mind how Ksp is calculated. Wikipedia gives a solubility of 3.18 g/L for TlCl. Anyways, I expect Tl compounds to be in the same price range as Ag. Add that to the limited availability and the toxicity and there's no real advantage to using Tl.

Quote: Originally posted by Vargouille

So, for quantitative purposes, AgNO3 is probably your best bet.

I concur. There's a reason titration with AgNO3 is the standard method despite the price and photosensitivity.

Quote: Originally posted by Vargouille

EDIT: Color me impressed. Sigma-Aldrich sells >97% Hg2Cl2-2H2O for $73/50g and >99.0% AgNO3 for $97/25g. So, you might be able to use the Hg2(NO3)2 for your quantitative purposes, but the toxicity is still an issue. Purity may be an issue as well, but that's nothing a recrystallization couldn't fix, eh?

Hg⁺ isn't very stable in solution, mercury(I) nitrate reacts slowly with water to give a precipitate. Also, salts of Hg⁺ disproportionate to Hg and Hg²⁺.

Swede: My recommendation to you is to use silver and recycle it. The precipitate can be filtered off and recycled easily. If that's really out of the question you should look into some spectrophotometric methods.

ScienceHideout - 16-9-2012 at 15:46

Can't you titrate with dichlorofluorescein somehow?

Vargouille - 16-9-2012 at 16:23

I would classify a solubility of 3.18g/L as (in the vernacular) "fairly insoluble". Perhaps sparingly soluble would have been the better turn of phrase, and it's a tenuous point to argue, I'll admit.

From what I can learn, the use of dichlorofluoroscein is used in the Fajans method of determining the concentration of chloride as a way to more easily see the end point. It doesn't eliminate the use of AgNO3, unless there is another usage I am overlooking.

Regarding the mercurous salt, could it not be used in a manner analogous to the Fajans method with a mercurous nitrate solution prepared immediately beforehand? There would be some standardization to be done, granted, and the waste dealt with as necessary, but the difference in solubility between silver chloride and mercurous chloride would aid in more accurate chloride determination.

EDIT: I found a reference for it, as the article Titration of Chloride and Bromide with Mercurous Nitrate Using Brom Phenol Blue as Adsorption Indicator, and it mentions other articles with mercurous nitrate.

[Edited on 17-9-2012 by Vargouille]

ScienceHideout - 16-9-2012 at 16:53

Hmmmm... You can react a crapload of lead nitrate with a sample of the stuff that has the chloride ion, making sure that the chloride is the limiting reagent, and then weigh the precipitate... do stoiciometry...

EDIT: Anything that has an insoluble chloride would work: You can use lead acetate, maybe...

[Edited on 17-9-2012 by ScienceHideout]

Lambda-Eyde - 16-9-2012 at 17:35

Quote: Originally posted by Vargouille

Regarding the mercurous salt, could it not be used in a manner analogous to the Fajans method with a mercurous nitrate solution prepared immediately beforehand? There would be some standardization to be done, granted, and the waste dealt with as necessary, but the difference in solubility between silver chloride and mercurous chloride would aid in more accurate chloride determination.

EDIT: I found a reference for it, as the article Titration of Chloride and Bromide with Mercurous Nitrate Using Brom Phenol Blue as Adsorption Indicator, and it mentions other articles with mercurous nitrate.

Thanks for proving me wrong. It goes to show how you can't always trust your own intuition.

I would still say that this doesn't bring Swede closer to his goal, though. Taking molecular weight into consideration, mercurous nitrate isn't going to be much cheaper. Add the toxicity and the concerns of disposing/recycling it afterwards, and I would in Swede's place gladly pay a little extra for the silver.

BromicAcid - 16-9-2012 at 19:50

Maybe a bit obvious but a chloride ion selective electrode would fit the bill.

AJKOER - 16-9-2012 at 22:19

More hard work, but perhaps a less costly route that directly verifies the presence of Chlorine.

1. Prepare fresh pure dilute HOCl (from say distilling vinegar and dilute NaOCl). Keep cool and in the dark for prompt use as decomposition and disproportionation will create HCl. To test concentration add known enough of HCl and note volume of Cl2.

2. Add an excess of Hypochlorous acid to your unknown chloride source.

3. Add a strong cheap acid or acid salt (like NaHSO4) in excess (but not HCl).

4. Note the formation of any Chlorine gas. Collect the Cl2. Based on the volume and solubility of Cl2 in the medium you used, one could calculate the implied chloride concentration if any. Per the reaction:

HOCl + H(+) + Cl(-) <--> Cl2 (g) + H2O

each mole (22.4 liters) of the observed Chlorine gas corresponds to a single mole of chloride. This equation also suggests, as is known, that Chlorine is largely insoluble in acidic solutions, so collect the Cl2 over a low pH water solution. Note, one can replace HOCl with a hypochlorite salt of known concentration. Also, one could elect to add some CaCl2 to the Chlorine water and cool to around zero degrees Centigrade and harvest Chlorine hydrate, formula Cl2 with 6.01 to 7.63 H2O depending on preparation, see https://docs.google.com/viewer?a=v&q=cache:vvyh40wfT0oJ:... .

[EDIT] Note, any soluble Silver salt added to a solution containing any dissolved CO2 will readily form AgCO3, an insoluble light yellow colored salt. However, the presence of CO2 is not an issue for the hypochlorite based tested described above. Both the HOCl and AgNO3 are subject to decomposition in the presence of light.

[Edited on 17-9-2012 by AJKOER]

Vargouille - 17-9-2012 at 01:53

Quote: Originally posted by Lambda-Eyde

I would still say that this doesn't bring Swede closer to his goal, though. Taking molecular weight into consideration, mercurous nitrate isn't going to be much cheaper. Add the toxicity and the concerns of disposing/recycling it afterwards, and I would in Swede's place gladly pay a little extra for the silver.

Oh my, I have made quite the error. Silver nitrate is actually cheaper on a basis of $/mole (at least from Sigma-Aldrich), and becomes more so the larger the order. Mercurous nitrate gets close, but not close enough to justify it unless you're more concerned with extreme accuracy.

So yes, unless you are preforming analytical chemistry to such a degree that the extra accuracy of mercurous nitrate is necessary, go with the silver nitrate.

The issue I have with the hypochlorite is that the other constituents of the chloride solutions are, as of writing, unknown. To use this method, you would have to account for those reactions, as well as the reactions with the chlorine released. It could work, but the application seems restricted.

AJKOER - 17-9-2012 at 05:21

Quote: Originally posted by Vargouille

The issue I have with the hypochlorite is that the other constituents of the chloride solutions are, as of writing, unknown. To use this method, you would have to account for those reactions, as well as the reactions with the chlorine released. It could work, but the application seems restricted.

With the exception of soluble flourides, nitrates, chlorates and perchlorates, all other Silver salts that could be produced are insoluble salts in the presence of acidic AgNO3. So how is not knowing the other unknown constituents a positive for AgNO3 over HOCl?

Example, any sulfides presence would form some AgS with acidic AgNO3. Yes, there is a problem with HOCl also, as it could react as follows with the H2S:

H2S + HOCl (in excess) --> HCl + H2O + S (s)

which would form Cl2 from the HCl even in the absence of any chloride. Note, however, there is a clue by the formation of a Sulfur suspension.

Another example, a Sulfate is present but not chloride. Adding AgNO3 could still form some insoluble colorless Ag2SO4, but the HOCl method is not effected.

[Edited on 17-9-2012 by AJKOER]

Hexavalent - 17-9-2012 at 07:59

It's not perfect, but you might want to look into the Beilstein test for your halide analyses....

http://en.wikipedia.org/wiki/Beilstein_test

Lambda-Eyde - 17-9-2012 at 08:16

Quote: Originally posted by Vargouille

Do keep in mind that one mole of mercurous nitrate precipitates two moles of chloride.

Quote: Originally posted by Hexavalent

It's not perfect, but you might want to look into the Beilstein test for your halide analyses....

http://en.wikipedia.org/wiki/Beilstein_test

That's a qualitative, not a quantitative test.

And the HOCl method doesn't exactly look like a quantitative procedure, either...

Vargouille - 17-9-2012 at 10:57

Then it depends on how you calculate how much it costs. I was calculating on moles of the nitrate. By moles of chloride reacted, however, the cost of Hg2(NO3)2 is between $409.69 (at 50g) and $288.47 (at 250g), and the cost of AgNO3 is the same. So if the only concern is in this titration, mercury nitrate is a fair bit less expensive.

And, regarding the HOCl/AgNO3 viability concern, the presumption from the OP was that the solutions were chlorides, with little (if any) contamination from other anions. In that case, yes, it would have to be amended to remove sulfide, carbonate, hydroxide, what have you. In the same case, those contaminating salts would skew the results of the HOCl as well, as the usefulness of a redox titration is diminished if multiple components of the solution can react with the titrant.

The benefit of a AgNO3 titration is that, once competing anions are removed, the cations will rarely precipitate from solution as the nitrates. The concern with HOCl is that the presence of transition metal ions can catalyze decomposition (see JAWWA volume 103 number 6) and cause reactions with the HOCl (as with ferrates and permanganates). The JAWWA article even suggests that commercial samples of NaOCl may require extra purification to remove transition metal ions.

EDIT: Admittedly, there are some transition metal ions that will react with silver ions, but they are fewer in number than those that react with HOCl.

[Edited on 17-9-2012 by Vargouille]

AJKOER - 17-9-2012 at 12:11

Quote: Originally posted by Vargouille

And, regarding the HOCl/AgNO3 viability concern, the presumption from the OP was that the solutions were chlorides, with little (if any) contamination from other anions. In that case, yes, it would have to be amended to remove sulfide, carbonate, hydroxide, what have you. In the same case, those contaminating salts would skew the results of the HOCl as well, as the usefulness of a redox titration is diminished if multiple components of the solution can react with the titrant.

The benefit of a AgNO3 titration is that, once competing anions are removed, the cations will rarely precipitate from solution as the nitrates. The concern with HOCl is that the presence of transition metal ions can catalyze decomposition (see JAWWA volume 103 number 6) and cause reactions with the HOCl (as with ferrates and permanganates). The JAWWA article even suggests that commercial samples of NaOCl may require extra purification to remove transition metal ions.

EDIT: Admittedly, there are some transition metal ions that will react with silver ions, but they are fewer in number than those that react with HOCl.

As I made it very clear in my recommendation, start with fresh distilled dilute HOCl (no transition metal ions). I am, while recognizing your point, assuming that the various other decomposition reactions associated with the HOCl (from light, lower pH, raise in temperature, disproportionation, transition metal presence, etc.) occurring in the test solution would proceed relatively slowly. So rapidly collect any Cl2 formed and disconnect the Chlorine generator. This is in contrast with AgNO3 where, until the very completion of filtering and washing, the secondary reactions (including reacting with any CO2 in the air) could continue.

AJKOER - 20-9-2012 at 05:29

Quote: Originally posted by Vargouille

EDIT: Admittedly, there are some transition metal ions that will react with silver ions, but they are fewer in number than those that react with HOCl.

Well, here is an extract from Wiki on a troubling ligand that the AgCl interacts with, namely Chloride (yes, the same that we are testing for), but primarily (I can only guess) at higher concentrations as concentrated HCl apparently does dissolve AgCl. To quote (source http://en.wikipedia.org/wiki/AgCl ):

"AgCl dissolves in solutions containing ligands such as chloride, cyanide, triphenylphosphine, thiosulfate, thiocyanate and ammonia. Silver chloride reacts with these ligands according to the following illustrative equations:

AgCl(s) + Cl−(aq) → AgCl2−(aq)
AgCl(s) + 2S2O32−(aq) → [Ag(S2O3)2]3−(aq) + Cl−(aq)
AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Cl−(aq) "

and, on AgCl solubility: "soluble in NH3, conc. HCl, conc. H2SO4, alkali cyanide, NH4CO3?, KBr, Na2S2O3; "

Of course to mitigate this potential problem, one can dilute the test solution and/or use an excess of the Silver nitrate. One probably should add the test solution slowly with stirring to the aqueous AgNO3 to avoid potential local high concentrations of Chloride or other problematic ligands.

[Edited on 20-9-2012 by AJKOER]

Swede - 20-9-2012 at 06:28

I appreciate all the thought that has gone into this thread. An ion-selective probe is expensive, but in the end, may be the most cost-effective choice. Otherwise, silver nitrate seems to be the gold standard.

Vargouille - 20-9-2012 at 11:26

How exactly does one prepare a solution of "fresh, distilled, dilute HOCl"? HOCl only exists in solution, and doesn't "distill" as you refer to it. But you meant the acidification of a sodium hypochlorite solution, I take it. Therein lies the contamination. Transition metal ions are difficult to remove completely from commercial NaOCl. Typically, NaOCl is available as solution, and only as reagent grade as a concentration range, which means that there is likely to be chloride and chlorate in solution. So acidification will shift equilibrium to form Cl2. Then you have to account for that. But there are more stable salts of hypochlorite, typically Ca(OCl)2. Except that only the technical grade is available from Sigma-Aldritch to anyone who can't tell the difference between the Minshu-tou and the Jimin-tou. Perhaps reagent grade Ca(OCl)2 is available elsewhere.

Of course, you could make Cl2O, and dissolve it in water to form HOCl, but that has the amazingly entertaining conditions of being most easily produced by a reaction between Cl2 and a mercury salt, and hydrolyzing in an equilibrium. And there is the issue of the quantitative collection of the Cl2. Cl2 is partially soluble in water and hydrolyzes. The only way to get it to completely exit the solution is to boil it. Then you have the problem of all of the equilibriums. HOCl will reverse into Cl2O and H2O at elevated temperatures. With chloride, it will also reverse into Cl2. Then, because it's acidified, HCl will exit the solution at high temperatures. How do you determine how much of all the gas you collect at the end belongs to each of these constituents? Far as I can tell, long hours in the library looking up equilibria.

As for the compounds that solubilize AgCl, concentrated H2SO4, HCl, and NH3 are admittedly troublesome for an AgNO3 titration. Thus, they can be tested for and isolated for testing with probes (CaCl2 for H2SO4, HCl and NH3 by smell, or by addition of CuCl2 to form the relevant complexes). NH3 is a problem for the HOCl method as well, rapidly increasing the pH of the HOCl solution, leading to a whole host of problems. AgCl2- is an issue with excess chloride, and is negated once the excess is removed. Thiosulfate can be tested for with iodine or bromine, and if the problem arises, set aside for use with the ion probe. It also reacts with Cl2 and HOCl in the HOCl method. The alkali cyanide is a problem in either case (solubilizes AgCl; releases HCN on contact with acidified HOCl). The bromide is a problem in both cases: precipitates as AgBr, the solubilizing of AgCl can be rectified with more AgNO3; and will be oxidized by Cl2 and HOCl.

[Edited on 20-9-2012 by Vargouille]

AJKOER - 21-9-2012 at 07:29

Quote: Originally posted by Vargouille

How exactly does one prepare a solution of "fresh, distilled, dilute HOCl"? HOCl only exists in solution, and doesn't "distill" as you refer to it. But you meant the acidification of a sodium hypochlorite solution, I take it. Therein lies the contamination. Transition metal ions are difficult to remove completely from commercial NaOCl. Typically, NaOCl is available as solution, and only as reagent grade as a concentration range, which means that there is likely to be chloride and chlorate in solution. So acidification will shift equilibrium to form Cl2. Then you have to account for that. But there are more stable salts of hypochlorite, typically Ca(OCl)2. Except that only the technical grade is available from Sigma-Aldritch to anyone who can't tell the difference between the Minshu-tou and the Jimin-tou. Perhaps reagent grade Ca(OCl)2 is available elsewhere.

Of course, you could make Cl2O, and dissolve it in water to form HOCl, but that has the amazingly entertaining conditions of being most easily produced by a reaction between Cl2 and a mercury salt, and hydrolyzing in an equilibrium. And there is the issue of the quantitative collection of the Cl2. Cl2 is partially soluble in water and hydrolyzes. The only way to get it to completely exit the solution is to boil it. Then you have the problem of all of the equilibriums. HOCl will reverse into Cl2O and H2O at elevated temperatures. With chloride, it will also reverse into Cl2. Then, because it's acidified, HCl will exit the solution at high temperatures. How do you determine how much of all the gas you collect at the end belongs to each of these constituents? Far as I can tell, long hours in the library looking up equilibria.

As for the compounds that solubilize AgCl, concentrated H2SO4, HCl, and NH3 are admittedly troublesome for an AgNO3 titration. Thus, they can be tested for and isolated for testing with probes (CaCl2 for H2SO4, HCl and NH3 by smell, or by addition of CuCl2 to form the relevant complexes). NH3 is a problem for the HOCl method as well, rapidly increasing the pH of the HOCl solution, leading to a whole host of problems. AgCl2- is an issue with excess chloride, and is negated once the excess is removed. Thiosulfate can be tested for with iodine or bromine, and if the problem arises, set aside for use with the ion probe. It also reacts with Cl2 and HOCl in the HOCl method. The alkali cyanide is a problem in either case (solubilizes AgCl; releases HCN on contact with acidified HOCl). The bromide is a problem in both cases: precipitates as AgBr, the solubilizing of AgCl can be rectified with more AgNO3; and will be oxidized by Cl2 and HOCl.

Vargouille: Here is one simple inexpensive manner to prepare a largely chloride free HOCl: acidify a hypochlorite with either a very dilute mineral acid or a weak acid (Boric, acetic,...) and distill. To quote from Watts' Dictionary of Chemistry, Volume 2, page 16:

"A dilute solution of HCl0 may be distilled with partial decomposition, the distillate is richer in HCl0; Gay-Lussao found that, on distilling a dilute solution to one-half, the distillate contained five-sixths of the total HClO (C. R. 14, 927)"

Another reference suggests twice distilled to further concentrate as needed. My understanding of why this concentrating effect is occurring is that the gaseous anhydride of HOCl (Dichlorine Monoxide, Cl2O) along with the volatile HOCl is boiled off 1st. For the current application, freshly prepared dilute HOCl kept cool and out of sunlight should suffice.
_______________________________

Now, consider a reaction with a test solution containing no Chloride, but an Iodide. With AgNO3, this becomes a recommended method for the preparation of AgI. Per Wiki (http://en.wikipedia.org/wiki/AgI ):

"Silver iodide is prepared by reaction of an iodide solution (e.g., potassium iodide) with a solution of silver ions (e.g., silver nitrate). A yellowish solid quickly precipitates."

Now, an acidified Iodide solution forms some aqueous HI. Upon slow addition of HOCl with stirring, one expected reaction:

2 HI + HOCl --> H2O + I2 + HCl

So a violet colored Iodine water solution should form. With more HOCl, Chlorine water (Cl2 gas) could be formed:

HCl + HOCl --> Cl2 + H2O

In the case of Bromide, similar reactions occurs with AgNO3 forming a soft, pale-yellow, water insoluble salt (AgBr) except that Silver bromide is unusual sensitivity to light.

I think the reactions with HOCl, certainly in the cases described, are more informative and perhaps, less prone to error, if you confuses the pale-yellow Silver salts for AgCl (perhaps more likely in the presence of some chloride).

Your comment on the dangerous reaction with a hypochlorite and a cyanide is questionable. In point of fact, NaOCl is commercially employed as a scrubber for cyanide (see, for example, http://www.monroeenvironmental.com/air-and-gas-cleaning-syst... ). A reaction, implied by this source, is:

NaCN + NaOCl --> NaCNO + NaCl

where cyanates are viewed as less toxic can be removed by the blowdown process. However, I do not believe that the following reaction with HOCl actually occurs:

HCN + HOCl --?--> HCNO + HCl

and even if does occur, to quote Wiki: "Isocyanic acid hydrolyses to carbon dioxide and ammonia:
HNCO + H2O → CO2 + NH3"

[Edited on 21-9-2012 by AJKOER]

Vargouille - 21-9-2012 at 11:50

Interesting. Presumably the equilibrium constant for 2HOCl -> H2O + Cl2O is greater than that for 3HOCl -> HClO3 + 2HCl. I am curious as to whether the HOCl is actually volatile. My understanding of the situation is that HOCl only exists in solution, and the characteristic smell of bleach solutions is due to it forming the gaseous anhydride or chlorine gas. In any case, I see no issue with forming the pure solution immediately beforehand, decomposition aside.

As for the other silver halides, note this picture of them. From left to right, it goes AgI, AgBr, AgCl, and AgF. Considering that, to my knowledge, titrations are typically preformed with a white or black backdrop as required, the difference between AgI and the others is fairly simply distinguished. To my eyes, AgBr has a slight dunnish cast to it, perhaps due to photodegradation, while AgCl is more close to pure white. If other halides are suspected, then it is a simple matter of preparing suspensions of the relevant halides beforehand, and if it comes up, reverting to an ion probe.

The question, however, of the cost and limitations of such an ion probe did come up. Surely it isn't omnipotent. A cursory search gave this quote:

Quote:

Because of the much greater solubility of AgCl compared to AgI, the Chloride electrode will be irreversibly damaged if immersed in solutions containing significant numbers of Iodide ions. Also note that all poly-crystalline membranes contain Silver Sulphide and thus will not give reliable readings if more than a trace of Ag or S ions are present in the solution. There is also a high interference from Bromide and Cyanide ions and the Chloride electrode will only give reliable results if these ions are absent, or only present in insignificant amounts compared to the Chloride ion.

If samples are likely to contain significant quantities of these ions, then their effect may be reduced by mixing samples and standards 1:1 with a sodium bromate buffer. This is made by dissolving 15.1 g sodium bromate in 800 ml water and addding 75ml of concentrated nitric acid. This must be stirred well and diluted to 1000 ml with water. Note that this is a strong oxidising solution which should be handled carefully and prepared and used in a well ventilated area since it may liberate Bromine gas. This buffer should remove up to 1000ppm of Bromide or Iodide and 500ppm Sulphide. Small quantities of Cyanide should also be oxidised - but this acid solution must not be added to samples with significant cyanide content because of the danger of liberating lethal HCN gas.

So even these ion probes use a silver salt. Unfortunately, they are not reliable if iodide, bromide, sulfide, or cyanide ions are in solution. The site does, fortunately, give the solution (except for the cyanide situation) in the form of a solution of bromic and nitric acids. The specific definition of what constitutes a "significant amount" of these anions is a bit of a mystery to me. Perhaps you can find a source to clarify. Prices on other sites range from $179 from Vernier for a student-quality probe, and $715 from Hach for what I presume is a professional-grade probe.

Nico 2000 Chloride ISE Technical Specifications