Dissolved Oxygen in Water Above Atm Pressures 4

[dberetta]

I am looking for data on the amount of oxygen that can be dissolved in water at pressures in the 30 to 120 psig range and room temperature. Are there any sources for this information?
Thanks

 

[btrueblood]

Look up Charles' Law in a good chem. handbook or textbook: "The concentration of a gas dissolved in a liquid at any given temperature is directly proportional to the partial pressure of the gas on the solution."

 

[shahyar]

The law is "Henry's Law", isn't it?

 

[mbeychok]

dberetta:

Henry's Law can be expressed as follows:

x = P/H

where:
x = mols of dissolved gas per mol of solution = mol fraction
P = absolute partial pressure of the gas above the solution, in atm
H = Henry's Law constant, in atm/(mol fraction)
1 atm = 1 atmosphere = 14.7 psia

The 6th Edition of "Perry's Chemical Engineers' Handbook" gives these values for the Henry's Law Constant of oxygen in water:

H = 36,400 atm/(mol fraction) at 15 deg C
H = 40,100 atm/(mol fraction) at 20 deg C
H = 43,800 atm/(mol fraction) at 25 deg C

I will assume that your pressure range of 30 to 120 psig refers to the pressure of your oxygen (rather than air pressure). Thus, as an example, for say 20 deg C and an oxygen pressure of 50 psig (i.e., 64.7 psia):

Oxygen pressure = 64.7/14.7 = 4.4 atmospheres
x = 4.4/40,100 = 0.00011 mol fraction of oxygen in water

If I assume that your pressure range of 30-120 psig refers to air pressure, then for 20 deg C and air pressure of 50 psig:

Air is approximately 20% oxygen and 80% nitrogen.
Thus, oxygen pressure = 0.20(64.7)/14.7 = 0.88 atmospheres
x = 0.88/40,100 = 0.000022 mol fraction of oxygen in water

One might argue that the partial pressure of the water vapor should be subtracted from the oxygen pressures above but, at 20 deg C, the vapor pressure of water is relatively insignificant compared to the above oxygen pressures.

I assume that you know how to convert the above answers in mol fractions to whatever concentration units you may desire.

Milton Beychok
(Contact me at www.air-dispersion.com)

 

[25362]

Shayar is right. See:

Boyle's -17th century- Law: V [ε] 1/P (const. T, n)
Charles's -18th century- Law: V [ε] T (const. P, n)
Avogadro's -19th century- principle: V [ε] n (const. P, T)
Ideal gas law, combines the above into: V [ε] n.T/P

Jaccques Charles and Joseph-Louis Gay-Lussac, both French scientists interested in balloons, made measurements of how the temperature of a gas affected pressure, volume and density. Charles built the first hydrogen balloon, which became known as the charli[è]re. Gay-Lussac established a world altitude record of 23,018 ft in 1804.

The English chemist William Henry, in 1801, established that the solubility of a gas in a liquid is proportional to its partial presure above the liquid. The law is normally written

Solubility = k_H [×] partial pressure​

The value of k_H, mol/(L.atm), at 20 deg C for gases in water are:

air 7.9 [×] 10^-4
oxygen 1.3 [×] 10^-3
CO₂ 2.3 [×] 10^-2
nitrogen 7.0 [×] 10^-4
hydrogen 8.5 [×] 10^-4

As an illustration, the partial pressure of oxygen at sea level is 0.21 atm, thus its solubility in water at 20^oC would be

0.21 atm [×] 1.3 [×] 10^-3 mol/(L.atm) = 2.7 [×] 10^-4 mol/L​

 

[dberetta]

Thank you all. It is rather embarrassing to realize that I forgot something like that. (Haven't done much chemical engineering since 1980)
Milton: I had dug out my (5th edition) of Perry's and couldn't find any references before I posted. I just looked again under constants, physical properties of pure substances, Henry's Law etc and couldn't find anything. I don't know if the 5th edition is missing something that the 6th has or ??? I am going to look one more time.
My application is feeding air into (drinking)water in a pipe under pressure via a venturi to oxidize H2S.
Best Regards
David Beretta

 

[25362]

David, since this is a slow chemical reaction it may be that you need much more air (oxygen) than estimated by stoichiometry, and some kind of filter to catch the elemental sulfur formed in the oxidation.

 

[dberetta]

25362: Thank you. We are doing a joint venture/partnership with a company that has a bit of experience installing aeration systems but is lacking in theoretical experience that might help them know what to expect as they try tougher applications. Your help and advice are appreciated. I had already tried to explain the need for an excess of O2 over stoichiometry so it is nice to have you also say so.
Best Regards

 

[btrueblood]

Whoops! I looked it up as I was writing the post yesterday, and even caught my mistake, but then didn't re-read the post before sending it on. Sorry about that.

 

[mbeychok]

dberetta:

In the Sixth Edition of Perry's, the Henry's Law constants are in Chapter 3, "Physical and Chemical Data" in a section labeled "Solubilities".

Milton Beychok
(Contact me at www.air-dispersion.com)

 

[mbeychok]

25362:

Just as a point of interest only:

k_H of 1.3 X 10^-3 mol/Liter-atm = H of 42,735 atm/mol fraction

which is pretty much the same as the value of H = 40,100 as given by Perry's for 20 deg C.

Milton Beychok
(Contact me at www.air-dispersion.com)

 

[Montemayor]

I can offer very little knowledge on Henry’s Law – over that expounded by Milton Beychok and 25362 but can offer my stars for their contribution. I’ve had both tasteful and distasteful experiences with Henry’s Law and I feel I’m still learning. It isn’t as simple as I would like it to be.

I have some thoughts on the subject for Milton and 25362 that are related to the theme discussed. My experience with Henry is that his Law is not exact and the properties of the dissolved solute, including its escaping tendency, change as the concentration changes and so K changes. This Law is credible and useful only in very dilute solutions. Up to a pressure of 1 atm, it holds within 1 to 5% accuracy with many gases. Above that pressure, it starts to generate erroneous design data. I don't have any data on the dilution limitations and I would appreciate anyone's experience, information, or comments on this point.

The mighty Dalton (of partial pressure legend) showed that the solubility of the individual gases in a mixture of gases is directly proportional to their partial pressures, the solubility of each gas being nearly independent of the presence of the others. Here, we find Dalton being the explicit, practical, individual that he was. Real world solubility applications rarely take place with only one, solitary pure gas in question – and the practical engineering question would be: does one gas affect the solubility of the other(s)? The solubility of oxygen in water is nearly twice as great as that of nitrogen; and since the solubility of one gas is unaffected by the presence of the other, the dissolved air in considerably richer in oxygen than the air above water. Here, we have a Unit Operation within another Unit Operation! In carrying out a solubility operation, we have also caused an “enrichment” of the dissolved air. I found this to be unique.

I have failed to find a credible, theoretical explanation of the solubility of inert gases in liquids. This subject was on the research agendas when I graduated, and I have found no reasonable progress made in this field in the last 40 years. Does anyone have information on this? Additionally, the solubility of any gas in liquids is usually decreased by the addition of other solutes – particularly electrolytes. I understand that the relative decrease in solubility is the same for different gases. Does anyone have recent information on this?

This subject of solubility and Henry’s Law is an interesting and practical one. We engineers often overlook its effects and causes when running process simulations and subsequently in the design of process equipment. For example, would feeding pure oxygen (or enriched air) be more productive or efficient in dberetta’s practical application? In light of what Dalton found, one would say no; but perhaps the total elimination of nitrogen’s partial pressure could offer rewarding effects – with the usual caution and concerns for handling pure oxygen.

 

[mbeychok]

Art:

You are correct, of course, in that Henry's Law applies only to very dilute solutions. My above calculation of 0.000022 mols of oxygen per mol of water (at 20 deg C and 50 psig air pressure) amounts to 39 ppm by weight ... which is indeed quite dilute.

I don't recall having ever seen any data as to the effect of solute gas pressure on the value of the Henry's Law constant ... but that doesn't mean that there is no such data, it just means that I haven't seen any.

Milton Beychok
(Contact me at www.air-dispersion.com)

 

[DickRussell]

If one were to look at a system of dissolved gases in water from the viewpoint of nonideality in the liquid phase, some nonideality would be expected, due to the dissimilarity in the species. But at exceedingly dilute solutions, the solvent (water in this case) is nearly pure, so that its activity coefficient is essentially 1.00, whereas the activity coefficients of each dissolved gas would be just about at its infinite dilution value. Even a ten-fold increase in such a low concentration still leaves a very low concentration, and so its activity coefficient would hardly change. Henry's law simply reflects this, but only for low concentrations.

At some point, the concentrations would become enough removed from "exceedingly low" that the activity coefficients would move significantly from their infinite dilution values and the individual dissolved gas concentrations would start to interfere with each other. Of course, at such high pressures needed to force such concentrations, vapor phase nonideality also would have to be considered, using an EOS. However, I won't claim to know just which combinations of liquid and vapor phase models would work well for such conditions. Fortunately, nearly all the time we deal with conditions for which the lumping of effects into Henry's law works well.

 

[siretb]

I have not found data for O2, but the book from Landolt & Börnstein states that some comparison can be made with N2 or He.
The Henry's constant, at a given temperature decreases slightly when P increases from 0 to 1000 bar

The gamma coefficient mols/kilogram/bar, for N2
at 40°C 5.5e-4 mol/kg/bar @ 1 bar
at 40°C 5.0e-4 mol/kg/bar @ 50 bar
at 40°C 3.8e-4 mol/kg/bar @ 400 bar

The absolute figures are probably correct to +-10%, what is intersting is the trend.
Hope this helps

 

[25362]

Henry's law (HL) proportionality constant can be expressed differently depending on the selected concentration units.

HL is applicable, as explained by Montemayor, when there is no reaction between the gas and the solvent, the solubility is sparing, the gas doesn't ionize in water, and the gas acts reasonably as an "ideal" gas.

Thus, when the gas partial presure is P and the concentration is x, the k_x should be defined as

k_x = lim (P/x)​

x[→]0

The greater solubility of oxygen (over nitrogen) in water is explained by a dipole/induced-dipole effect, which makes HL's general applicability even more narrow.
Non-polar propane and butane solubilities in water follow HL quite well although k_x tends to lower a bit at higher pressures as the gases deviate from ideality, as shown by siretb on nitrogen. HCl, of course, doesn't obey it at any pressure because of ionization.

There are many coefficients for gas solubility in liquids, most of them refer to pressures around 1 bar: the Bunsen coefficient, the Kuenen coefficient, the absorption coefficient, and the Ostwald coefficient. The last one is independent of pressure under the assumption that an ideal gas volume is inversely proportional to pressure, thus the dissolved gas volume doesn't change with an increase in pressure.

The solubility of gases in hydrocarbons is generally expressed by use of empirically-obtained Ostwald coefficients as determined by ASTM D2779-92(2002).

These coefficients are defined as the ratio of the volume of dissolved gas to the volume of the solvent liquid at the test temperature and pressure. A pressure proportionality parameter should be used when intending to express them in a mass basis.

One discipline very interested in gas solubility in hydrocarbons and synthetic chemicals I know of is Engineering Tribology. Books specializing on this subject may bring more data on gas solubilities at various pressure levels.

 

[red8648jc]

Hi, I am new to this forum. I found the above discussions very useful.

Does anyone know a good source for Henry's constants of various gases (oxygen, nitrogen, hydrogen, etc) in ORGANIC SOLVENTS? Only low pressure info is needed. Thanks.

 

[hacksaw]

red,

you should start a new query to keep the responses specific to your question