Sciencemadness Discussion Board

Cl2 + H2O

Traveller - 4-2-2013 at 23:09

From what I understand, adding Cl2 to H2O would go like this:

Cl2 + H20 = HCl + HOCl

I understand that HOCl is a weak acid and may not alter the ph much but, HCl is a different matter.

If we started out with fairly pure water at a ph of 7, is there a way to determine how much the ph of the solution will drop for a given amount of Cl2 dissolved in it?

weiming1998 - 5-2-2013 at 03:15

The equation shows that 1 mole of Cl2 will dissolve in water to produce a mole of HClO and a mole of HCl. HClO is a very weak acid, so we can ignore its effects on the pH. HCl is a very strong acid, and will dissociate completely in water to H3O+. The concentration of H3O+ is basically the pH. 0.1M HCl has a pH of 1, 0.01M HCl a pH of 2, and so on. Using the equation above, you can work out that 7.1g of Cl2 dissolved in 1L of water gives you a pH of 1.

woelen - 5-2-2013 at 03:27

Weiming1998, your answer is not correct. This would be the case if all of the chlorine were dissociated into HCl and HOCl. In reality, however, the equilibrium is very far to the side of Cl2+H2O. Only a very small part exists as HCl + HOCl. So, a solution of Cl2 in water is nearly neutral.

A solution of Cl2 in water, however, slowly decreases its pH. HOCl is unstable and slowly decomposes, especially at heat and in the presence of light. HOCl decomposes to HCl and O2 and also for a small fraction to HClO3. So, if Cl2 is dissolved in water, then over the days its pH slowly decreases and after several days there will be noticeable acidity in the water and most of the Cl2 is gone and the solution contains mainly HCl.

AJKOER - 5-2-2013 at 05:21

Quite intuitively pressure (per Henry's Law "the solubility of a gas is directly proportional to the partial pressure of gas over a liquid" well, at least, if we are not talking about extreme pressures where the relationship can become non-linear) and temperature (where the relationship is inverse for water, but proportional for organic compounds, and in both directions for noble gases, see http://chemwiki.ucdavis.edu/Physical_Chemistry/Physical_Prop... ) are also factors in setting the equilibrium.

And just when you think you have it understood, don't add any significant amounts of a highly soluble foreign compound. For example, the solubility of NH3 in water is negatively impacted by adding NaCl (there is actually a technique called salting-out).


[Edited on 5-2-2013 by AJKOER]

elementcollector1 - 5-2-2013 at 08:10

Quote: Originally posted by AJKOER  
For example, the solubility of NH3 in water is negatively impacted by adding NaCl (there is actually a technique called salting-out).

Useful to know when I'm trying to purify 5% ammonia.

Traveller - 5-2-2013 at 18:23

The reason I asked about Cl2 dissolved in water can be found in this excerpt from "Getting Gold: A practical treatise for prospectors, miners and students" by J.C.F. Johnson, F.G.S. (1898).

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The most scientific and perfect mode of gold extraction (when the conditions are favourable) is lixiviation by means of chlorine, potassium cyanide, or other aurous solvent, for by this means as much as 98 per cent of the gold contained in suitable ores can be converted into its mineral salt, and being dissolved in water, re-deposited in metallic form for smelting; but lode stuff containing much lime would not be suitable for chlorination, or the presence of a considerable proportion of such a metal as copper, particularly in metallic form, would be fatal to success, while cyanide of potassium will also attack metals other than gold, and hence discount the effect of this solvent.

The earlier practical applications of chlorine to gold extraction were known as Mears' and Plattner's processes, and consisted in placing the material to be operated on in vats with water, and introducing chlorine gas at the bottom, the mixture being allowed to stand for a number of hours, the minimum about twelve, the maximum forty-eight. The chlorinated water was then drawn off containing the gold in solution which was deposited as a brown powder by the addition of sulphate of iron.

Great improvements on this slow and imperfect method have been made of late years, among the earlier of which was that of Messrs. Newbery and Vautin. They placed the pulp with water in a gaslight revolving cylinder, into which the chlorine was introduced, and atmospheric air to a pressure of 60 lb. to the square inch was pumped in. The cylinder with its contents was revolved for two hours, then the charge was withdrawn and drained nearly dry by suction, the resultant liquid being slowly filtered through broken charcoal on which the chloride crystals were deposited, in appearance much like the bromo-chlorides of silver ore seen on some of the black manganic oxides of the Barrier silver mines. The charcoal, with its adhering chlorides, was conveyed to the smelting-house and the gold smelted into bars of extremely pure metal. Messrs. Newbery and Vautin claimed for their process decreased time for the operation with increased efficiency.

At Mount Morgan, when I visited that celebrated mine, they were using what might be termed a composite adaptation process. Their chlorination works, the largest in the world, were putting through 1500 tons per week. The ore as it came from the mine was fed automatically into Krom roller mills, and after being crushed and sifted to regulation gauge was delivered into trucks and conveyed to the roasting furnaces, and thence to cooling floors, from which it was conveyed to the chlorinating shed. Here were long rows of revolving barrels, on the Newbery-Vautin principle, but with this marked difference, that the pressure in the barrel was obtained from an excess of the gas itself, generated from a charge of chloride of lime and sulphuric acid. On leaving the barrels the pulp ran into settling vats, somewhat on the Plattner plan, and the clear liquid having been drained off was passed through a charcoal filter, as adopted by Newbery and Vautin. The manager, Mr. Wesley Hall, stated that he estimated cost per ton was not more than 30s., and he expected shortly to reduce that when he began making his own sulphuric acid. As he was obtaining over 4 oz. to the ton the process was paying very well, but it will be seen that the price would be prohibitive for poor ores unless they could be concentrated before calcination.

The Pollok process is a newer, and stated to be a cheaper mode of lixiviation by chlorine. It is the invention of Mr. J. H. Pollok, of Glasgow University, and a strong Company was formed to work it. With him the gas is produced by the admixture of bisulphate of sodium (instead of sulphuric acid, which is a very costly chemical to transport) and chloride of lime. Water is then pumped into a strong receptacle containing the material for treatment and powerful hydraulic pressure is applied. The effect is stated to be the rapid change of the metal into its salt, which is dissolved in the water and afterwards treated with sulphate of iron, and so made to resume its metallic form.

It appears, however, to me that there is no essential difference in the pressure brought to bear for the quickening of the process. In each case it is an air cushion, induced in the one process by the pumping in of air to a cylinder partly filled with water, and in the other by pumping in water to a cylinder partly filled with air.

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I am planning to replicate the solution they made by beginning with NaOCl and reducing the ph with minute additions of HCl.

Traveller - 6-2-2013 at 10:35

Also, does anyone know what a "gaslight revolving cylinder" is? It is mentioned in the third paragraph of the excerpt and I have been unsuccessful in finding any information on this antiquated item.

syslen - 8-2-2013 at 13:14

Also, you have to know that the reaction is an equilibrium, so it doesn't react completely. A lot of chlorine will escape or stay unreacted in solution.

Traveller - 12-2-2013 at 10:40

I'm not quite sure I understand what is meant by the word "equilibrium" in this instance. Could you explain this a little to me?

IrC - 12-2-2013 at 17:02

Quote: Originally posted by Traveller  
Also, does anyone know what a "gaslight revolving cylinder" is? It is mentioned in the third paragraph of the excerpt and I have been unsuccessful in finding any information on this antiquated item.


A hundred years ago (or thereabouts) buildings were lit with gas lighting. In large buildings such as government and military a room similar to a boiler room produced the coal gas. A large (10 foot diameter) cylinder was filled with coal and constantly rotated as it was heated by a coal fire to red heat. A pipe came out through a valve which was a big pool of liquid mercury metal which allowed coal gas to exit but kept air from sucking back in. The resulting gas was purified with chlorine and fed through piping throughout the building to all the gas lamps.

Unbelievable but true, imagine rooms where the lamps blew out filling rooms full of gas. Or at the least the Hg vapor which must have been coming out into every room and consider the hours officials spent in them. Talk about living dangerously. Government leaders must have been mad as hatters.

What I have not figured out is why this is still true when they have had electric lighting and mercury free air for so many decades. What is their excuse today I wonder.

http://books.google.com/books?id=hzPVEteQLRsC&pg=PA162&a...


It leaves the question: who the hell stood all day in that furnace room turning some kind of geared crank to keep the cylinder rotating. I wonder if someone thought of low power steam locomotion to keep it slowly rotating?


[Edited on 2-13-2013 by IrC]

Traveller - 12-2-2013 at 19:24

Thanks for the answer. Any idea what combustible (and non-combustible) gases were in this "gas"?

barley81 - 12-2-2013 at 19:34

If the coal was reacted with steam at red heat (wiki says this is the case), then the products include H2, CO, and CH4 with some CO2 and N2. If this gas leaks, it could kill you from CO poisoning.

IrC - 12-2-2013 at 20:22

I never researched the reaction further and the old book never mentioned the actual reaction nor the use of steam in it, good info barley81. I could not find much on it at all. I bet a patent search could find this monstrosity. I wonder if they diverted some of the steam going to the coal bed to keep it turning?

Must have been used in really big buildings the info states the 10 foot cylinder would light 16,000 lights. With all the gasses barley81 mentions how the hell were the rooms safe to stay in? I have seen pictures of old buildings and there would be burning lamps every so many feet apart. No exhaust other than into the room. The CO must have been reacted in the lamp flame to produce CO2? Even if true, we are are still talking about rooms full of suffocating gasses, and wherever lights had gone out even deadly gasses. Also I wonder in the big hotels in New York and San Francisco, how many never woke up to pay their bill? Did they have a large 'deadbeat' tenants problem?

Curiosity made me research this subject further. There are some old books out there which are very interesting. I understand more now about such events as San Fransisco burning down after the big quake at the turn of the century. Must have been many broken pipes carrying pressurized flammable gas in all the large buildings. Some of these coal gas plants were enormous. Just imagine the plumbing in large cities to light street lamps. Especially dangerous in cities in earthquake prone areas. I have to say I prefer electric lighting after looking closer at what daily 1890 life was like.



[Edited on 2-14-2013 by IrC]

Traveller - 25-2-2013 at 00:32

quote from woelen:

"A solution of Cl2 in water, however, slowly decreases its pH. HOCl is unstable and slowly decomposes, especially at heat and in the presence of light. HOCl decomposes to HCl and O2 and also for a small fraction to HClO3. So, if Cl2 is dissolved in water, then over the days its pH slowly decreases and after several days there will be noticeable acidity in the water and most of the Cl2 is gone and the solution contains mainly HCl."

Thank you for this answer, woelen. It explains quite a bit to me.

If the solution of Cl2 in water, which we will assume to be HOCl and HCl if the pH of the solution is 7.5, was immediately placed in a sealed vessel and air pumped in at 60 psi and maintained at that pressure, would this pressure retard the decomposition of HOCl to HCl and O2?

The only thing I can compare this to is the HOCl I create when I add, by injection, sodium hypochlorite to our town's drinking water. As long as the water is in our pipelines at 94 psi, the HOCl has a relatively long life. However, if someone pours our water into a glass and leaves it on a counter overnight, the free chlorine, which I assume to be HOCl, disappears almost completely. Is this what is happening?

woelen - 25-2-2013 at 09:09

I do not think that decomposition of hypochlorite and HOCl is retarded by higher pressure. I believe that also in the pipelines of the water supply network there is decomposition. But as there is a constant consumption of water, the water never stays in the pipelines for more than several days. In the cold and dark pipelines the hypochlorite decomposes at such a low rate, that it usually makes it to the end users when they tap some water from the network.

I can imagine, that if you keep all taps closed for several weeks in a certain house, that the first few minutes of water after opening a tap will contain nearly no hypochlorite.

Traveller - 25-2-2013 at 09:54

That being said, I am at a loss to explain this part of the procedure I quoted from the mining treatise several posts back:

"They placed the pulp with water in a gaslight revolving cylinder, into which the chlorine was introduced, and atmospheric air to a pressure of 60 lb. to the square inch was pumped in."

I know this procedure made a solution of HOCl and HCl in water but, what would be the purpose of the air in the cylinder at 60 psi? From what I have read, elevating pressure does not seem to make any difference to the solubility of Cl2 in water.

A similar thing was described in a paper from the 1890's that talked about preparing a "chlorine solution" by the electrolysis of salt water. I assume they were making NaOCl but it could have been Cl2 gas. Either way, it would ultimately have become a solution of HOCl. From this paper:

"In order to treat the ore more effectually with the chlorine solution it was advantageous to expel the air from the chlorinator. For this purpose the chlorinator was provided with a valve, so that the air contained in the chlorinator passed out as the chlorine solution passed in. the valve was closed immediately the air was expelled."

Two different processes but, would expelling air from a chamber full of liquid not give them almost the same thing as an air pocket in the chamber at 60 psi? And what purpose would this serve?

ScienceSquirrel - 25-2-2013 at 10:17

I suspect that chlorine loss from a glass is by the decomposition of hypochlorite to chloride and oxygen and by the loss of chlorine to the air.
This would depend on temperature, the presence of light, etc.
Chlorine is also lost in pipe systems but at a much slower rate.

woelen - 25-2-2013 at 13:47

Solubility of chlorine in water is affected by partial chlorine pressure, not by pressure alone. But at high pressure of course the partial chlorine pressure can be higher as well. If this indeed is the case in the process, then higher pressure does have a strong influence.

Traveller - 25-2-2013 at 16:52

Quote: Originally posted by ScienceSquirrel  
I suspect that chlorine loss from a glass is by the decomposition of hypochlorite to chloride and oxygen and by the loss of chlorine to the air.
This would depend on temperature, the presence of light, etc.
Chlorine is also lost in pipe systems but at a much slower rate.


The average pH of the water in our system, after chlorination with NaOCl, is about 7.5. At this pH, free chlorine exists in our water at about a 50/50 mix of HOCl/OCl. As HOCl is the more unstable of the two by far, I would think that the major decomposition of free chlorine in a glass would be the loss of oxygen from HOCl and the subsequent creation of HCl. As pH mandates the 50/50 shift at pH 7.5, OCl will convert to HOCl, as quickly as the HOCl is lost to HCl and O2, until all of the free chlorine has decomposed as HOCl.

I was taught that the decomposition of NaOCl to NaCl and NaClO3 takes much longer and often requires warmer temperatures.

In closed pipelines, we see only loss of free chlorine to chloramines, trihalomethanes and other chlorinated disinfection byproducts.

[Edited on 26-2-2013 by Traveller]

Traveller - 25-2-2013 at 17:23

Quote: Originally posted by woelen  
I do not think that decomposition of hypochlorite and HOCl is retarded by higher pressure. I believe that also in the pipelines of the water supply network there is decomposition. But as there is a constant consumption of water, the water never stays in the pipelines for more than several days. In the cold and dark pipelines the hypochlorite decomposes at such a low rate, that it usually makes it to the end users when they tap some water from the network.

I can imagine, that if you keep all taps closed for several weeks in a certain house, that the first few minutes of water after opening a tap will contain nearly no hypochlorite.


The real question here is, what route of decomposition does HOCl/OCl follow while under pressure in our pipelines?

In storage, the major route of decomposition (about 90%) for NaOCl is:

3 NaOCl = 2 NaCl + NaClO3

while the minor route of decomposition (10%) is:

2 NaOCl = 2 NaCl + O2

As Health Canada limits the amount of chlorate in drinking water, we are obligated to store our bleach in almost refrigerated conditions and to keep the jugs well sealed; supposedly to provide backpressure to limit the offgassing of O2.

As our town is only 400 people and quite spread out, with many deadend branchlines, it cannot be said that the water goes through our lines quickly and is immediately consumed within a day or two. Many deadends have the same water sitting in the 200 mm line for a couple of weeks. When the free chlorine residuals do finally disappear, total chlorine tests often show the chlorine to be tied up in disinfection byproducts such as chloramines, haloacetic acids and trihalomethanes.

Just as an experiment, I tested the free chlorine in the water at my kitchen tap at 5:00 PM (1700). I live on one of these deadends and, beside my neighbour, we are the only consumers on this line. After 15 minutes of contact time, this water tested, at the pumphouse just down the road, at 1.3 ppm free chlorine. At my tap, it just tested at .8 ppm free chlorine. I poured a glass of this water and put it in the fridge where it is nice and dark. Our groundwater is roughly 9° C. so the temperature in the fridge should be similar.

I will test it tonight to see what the free chlorine residual is.

[Edited on 27-2-2013 by Traveller]

AJKOER - 25-2-2013 at 19:46

Is your water particularly rich in Iron?

If so, for water rich in Iron (from Fe(HCO3)2 for example), a reaction could occur further decomposing the HOCl more rapidy. My speculation based on the known effect of O2 on Fe(HCO3)2:

4 Fe(HCO3)2 + 2 HOCl + 2 H2O --> 4 Fe(OH)3↓ + 8 CO2↑ + 2 HCl

This means that the reaction:

Cl2 + H2O <--> HOCl + HCl

moves more to the right and the disinfectant properties attributed to HOCl is reduced. This implies to me the need for more Cl2 to be added.

Another point, especially with the added Chlorine, is the formation of FeCl3:

Fe(OH)3 + 3 HCl <--> FeCl3 + 3 H2O

which has a noted ability to increase the 'activity' of even dilute acids (HCl, for example) present particularly in mineral rich water. Bottom line, for water rich in Fe and other minerals, I would also expect relatively more evidence of corrosive activity on pipes, etc.


[Edited on 26-2-2013 by AJKOER]

Traveller - 25-2-2013 at 19:54

One of our wells has a higher iron and manganese content. However, since we are disinfecting with NaOCl at the pumphouse and NaOCl is an oxidizer, the Fe and Mn get oxidized almost instantly and precipitate out of solution. We do not filter these oxides but they seem to collect in our main pipeline and get flushed out when the mainline gets flushed.

Traveller - 26-2-2013 at 07:43

Okay, it is 6:18 AM here, just over thirteen hours since I placed the glass of chlorinated water in the fridge. Thirteen hours ago, it tested at .8 ppm free chlorine in water coming from a pipe where water temperatures are typically 9° C.

The first noticeable thing about the glass of water were the large number of air bubbles clinging to the glass under the water (O2?). I then tested with my testing kit and the water in the glass tested at 0.0 ppm free chlorine. This same water, at this time of year, usually can sit stagnant, at 94 psi, in the deadend pipeline that feeds my house for a minimum of two weeks before the free chlorine is depleted to this point.

I am currently out of total chlorine testing reagents but I will make a point of obtaining some. With these, it is possible to see just how much free chlorine combines with other compounds to make chloramines, THM's, HAA's, etc. If tap water is tested for free chlorine and total chlorine, as it comes out of the tap, and, once the free chlorine is totally gone, tested again for total chlorines, it will be possible to establish just how much of the HOCl/OCl follows the decomposition route into HCl and O2.

[Edited on 26-2-2013 by Traveller]

[Edited on 26-2-2013 by Traveller]