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WARNING: This procedure involves hot, concentrated hydrochloric acid, which is highly corrosive. To use this website or any information contained within it, you must read and agree to the Terms of Use. If you choose to attempt this or any of the other experiments discussed on this web site, you do so entirely at your own risk.
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Purification of Hydrochloric Acid
with a simple micro-still
by Christian Thorsten
Abstract: A micro still for sub-boiling distillation of HCl was tried using only the most readily-available laboratory items. The goal was to clarify a sample of iron-contaminated HCl and make it at least suitable for qualitative analysis and general lab use. Titration of the distillate suggested a final concentration of 7.1 to 7.3 M, although higher concentrations seemed attainable.
Introduction:
Hydrochloric acid's importance in the laboratory is difficult to overstate.
HCl is undoubtedly one of the most useful reagents in all of laboratory science.
There is no substitute for it.
Sometimes the need arises to purify small amounts of HCl for qualitative
or quantitative analysis experiments. Distilling an acid at a temperature
just below its boiling point reduces spattering and aerosolization of metal
ions and other contaminants, leading to a highly pure distillate. Prior
literature has shown ( ref. needed
) that even ACS grade HCl can be improved in purity by sub-boiling
distillation. In the case of acid that's become grossly contaminated
from metal ions such as iron, sub-boiling distillation can make reagent-grade
or better acid from just one or two steps.
A prior method ( ref. needed
) suggests leaving a dish of impure HCl and a dish of distilled H2O
next to each other in a closed container; after standing for some length
of time, the dish of water will have absorbed enough HCl vapor to make it
half the concentration of the contaminated sample. For example, a contaminated
sample of 12 molar HCl and a dish of pure water will eventually both become
6M HCl.
The author of the present study wondered if there was a way to improve the
above method so that stronger HCl could be attained in a comparably simple
apparatus. The results were promising.
Materials & Methods:
All glassware consisted of borosilicate glass.
The starting batch of HCl was 10 to 12 M concentration and had a visible, yellow coloration (presumably Fe3+ impurities).
I. Distillation
A 10-mL micro beaker was placed into a 250-mL beaker to which had been added
about 20 mL of impure hydrochloric acid. The micro beaker was centered,
and the larger beaker was fitted with a 75-mm funnel that had been specially
prepared by melting off the funnel stem in a gas flame to produce a drip point
(figure 1). The funnel / beaker joint was sealed with Parafilm and
the funnel top filled with ice. The beaker was heated on a laboratory
hot plate on lowest setting, just hot enough to cause very minor bubbling
in the HCl.
.
II. Determination of Acid Strength
The purified acid was diluted 1:100 by using a volumetric pipette to transfer 1.00 mL (+/- 0.007 mL) to a 100 mL volumetric flask. The volume was then made up to 100.0 mL with distilled water. Initial concentration was presumed to be in the vicinity of 10 molar, so this dilution was expected to give acid of somewhere around 0.1 molar.
Anhydrous sodium carbonate was prepared by heating the decahydrate at 150-160 °C for 3 hours in an oven. The dried material was stored in a vacuum desiccator until cool. The laboratory was at 48% relative humidity; mass change due to hydration was expected to be minimal during the few minutes the Na2CO3 was exposed to the air.
A 0.10 molar sodium carbonate solution was prepared by dissolving 2.650 grams of anhydrous Na2CO3 (fwt. 105.99) in enough distilled H2O to make 250.0 mL of solution.
The indicator chosen was methyl orange, which happens to have its color midpoint at the second equivalence point of the HCl / sodium carbonate titration curve. This color change is at pH 3.7 and goes from yellow to orange.
5.00 mL (+/- 0.015 mL) of the 0.100 M sodium carbonate solution were put in a small beaker using a volumetric pipette. To this was added a carefully-measured amount of the unknown (ca. 0.10 M) HCl solution until a color change was noted. A high precision micro burette (capacity 5.00 mL) was used to deliver the solution; three fillings were necessary to provide the necessary amount of HCl, the equivalence point being reached about halfway through the third aliquot of 5 mL. We kept track of the exact amount required to reach color-change.
Localized color change began to appear around 7.40 mL of added HCl, becoming pronounced by 12.40 mL; this went away with swirling. The complete change-over of the solution from orange to red-orange happened at 13.73 mL.
Since there were calculated to be 5.00 x 10-4 moles sodium carbonate in the beaker, and since the methyl orange equivalence point was deemed to occur where there were 2 moles HCl per 1 mole Na2CO3, there must therefore have been 10.00 x 10-4 moles HCl added to the beaker.
0.00100 moles HCl in 13.73 mL comes out to 0.0728 moles per liter. Since the original HCl sample was diluted 1 to 100, the concentration of HCl in the original sample was 7.28 moles / liter.
In another trial, the titration was found to require 14.14 mL of HCl, for a concentration of 7.07 moles / liter.
Discussion:
The micro-apparatus for purifying HCl was found to give 7.1 to 7.3 M HCl.
The titration methods used in this experiment led to at least a 3% margin
of error; at least some of this can be attributed to the failure of
methyl orange indicator to have a sharp color-change. Time permitting,
the ideal way to do this experiment would be (1) to use methyl orange + xylene
cyanol (OR use a different base standard and indicator), and (2) to conduct
multiple trials to assess the relative deviation and the mean.
The molarity of HCl determined in the present study reflects the concentration
obtained from three, combined distillation runs. Since 6 M is strong
HCl and is a standard laboratory concentration, the product of the micro still
is more than strong enough for general lab use.
When 37% hydrochloric acid is heated, HCl gas escapes until the solution
reaches about 20.2%. This azeotropic solution is about 6 M in HCl and will
distill over in this same concentration. The final product of the micro
still is therefore a somewhat enriched version of this.
It is possible that procedural modifications can enrich the acid beyond
the 7.1-7.3 M that was obtained in this experiment. Since no vapor
escapes the still until the seal is unwrapped, it is possible that concentration
can be increased by:
(1) redesigning the apparatus to have less internal air space in order to reduce the amount of undissolved HCl gas,
(2) refluxing the HCl from the micro beaker onto the cold funnel for a longer time, making sure to keep fresh ice in the funnel, and
(3) cooling the entire, sealed apparatus in an ice bath for perhaps 24 hours to help HCl gas pass back into solution before opening the system to the outside air.
If no excess HCl remains outside the micro beaker, the concentration of
HCl inside the micro beaker can theoretically approach 12 M, lessened by
whatever amount of HCl gas remains undissolved or tied up in any droplets
that remain on the inside wall of the beaker. Since the cold funnel
sends liquid on a one-way trip into the micro beaker, prolonged operation
of the system should be able to get fairly close to 12 M final concentration.
10 M HCl at least should be within reach.
If the micro beaker's maximum capacity is 15 mL, there should be no more
than 15 mL of impure HCl added to the system at the outset. This is
important to ensure maximal concentration of the acid.
Methyl orange is a useful indicator for the sodium carbonate-HCl titration,
but the equivalence point is not very sharp. Mixing it with some xylene
cyanole FF indicator, as is sometimes done in analytical labs, provides a
narrower color transition, although the colors involved are not the same.
Because Ka values that appear in books were derived experimentally
at some point, it is impossible to calculate the equivalence-point pH of
a titration with 100% confidence. The Ka values in different
sources may vary by as much as ~20%, even for the same listed concentration
and approximate temperature. If we assume +/- 20% error in the calculated
hydrogen ion concentration, that translates to an uncertainty of about 1/10
of a pH unit.
In the present experiment, the value 4.3 x 10-7, obtained from
the CRC Handbook of Chemistry and Physics (1988), was used.
Anhydrous Na2CO3 is often designated and sold as a
primary base standard; however, this is a puzzling choice when one considers
that anhydrous Na2CO3 absorbs moisture from the air.
Hygroscopic compounds do not make good primary standards; their mass
and composition can change during a laboratory session. Furthermore,
using carbonate in a titration adds an extra layer of complexity due to buffering.
A much better primary base standard is barium hydroxide. Nevertheless,
passable results can be obtained with sodium carbonate if the lab is kept
at low humidity and the compound is handled in such a way as to keep air exposure
to a minimum.
This experiment has shown a simple, convenient method for purifying small
amounts of hydrochloric acid. Future studies may refine the method
to enhance the concentration of HCl. While no trace metal analysis was
performed on the distillate, it is safe to assume the vast majority of metal
ions are removed during sub-boiling distillation.
by Christian Thorsten
Abstract: A micro still for sub-boiling distillation of HCl was tried using only the most readily-available laboratory items. The goal was to clarify a sample of iron-contaminated HCl and make it at least suitable for qualitative analysis and general lab use. Titration of the distillate suggested a final concentration of 7.1 to 7.3 M, although higher concentrations seemed attainable.
Introduction:
Hydrochloric acid's importance in the laboratory is difficult to overstate.
HCl is undoubtedly one of the most useful reagents in all of laboratory science.
There is no substitute for it.
Sometimes the need arises to purify small amounts of HCl for qualitative
or quantitative analysis experiments. Distilling an acid at a temperature
just below its boiling point reduces spattering and aerosolization of metal
ions and other contaminants, leading to a highly pure distillate. Prior
literature has shown ( ref. needed
) that even ACS grade HCl can be improved in purity by sub-boiling
distillation. In the case of acid that's become grossly contaminated
from metal ions such as iron, sub-boiling distillation can make reagent-grade
or better acid from just one or two steps.
A prior method ( ref. needed
) suggests leaving a dish of impure HCl and a dish of distilled H2O
next to each other in a closed container; after standing for some length
of time, the dish of water will have absorbed enough HCl vapor to make it
half the concentration of the contaminated sample. For example, a contaminated
sample of 12 molar HCl and a dish of pure water will eventually both become
6M HCl.
The author of the present study wondered if there was a way to improve the
above method so that stronger HCl could be attained in a comparably simple
apparatus. The results were promising.Materials & Methods:
All glassware consisted of borosilicate glass.
The starting batch of HCl was 10 to 12 M concentration and had a visible, yellow coloration (presumably Fe3+ impurities).
I. Distillation
A 10-mL micro beaker was placed into a 250-mL beaker to which had been added
about 20 mL of impure hydrochloric acid. The micro beaker was centered,
and the larger beaker was fitted with a 75-mm funnel that had been specially
prepared by melting off the funnel stem in a gas flame to produce a drip point
(figure 1). The funnel / beaker joint was sealed with Parafilm and
the funnel top filled with ice. The beaker was heated on a laboratory
hot plate on lowest setting, just hot enough to cause very minor bubbling
in the HCl..
|
Figure 1. The apparatus. The
funnel stem has been melted off in a gas flame, leaving a small protrusion
from which the condensed HCl will drip into the micro beaker.
The impure HCl is placed outside the micro beaker.
The upper portion of the funnel is filled with ice when distillation
commences.
The joint between the funnel and the large beaker is
sealed with Parafilm or PTFE tape before distillation commences. |
|
Figure 2. Heating the apparatus on
a hot plate (lowest setting). If the liquid actually boils, the setting should be turned down a bit. |
|
Figure 3. Nearly all the HCl ends
up in the micro beaker, minus the contaminating metal ions.
The yellow (contaminated) HCl in the photo is actually
outside the micro beaker.
The starting volume of acid outside the micro beaker
should be no more than the maximum capacity of the micro beaker. In
the photo, a little too much starting HCl was used. |
II. Determination of Acid Strength
The purified acid was diluted 1:100 by using a volumetric pipette to transfer 1.00 mL (+/- 0.007 mL) to a 100 mL volumetric flask. The volume was then made up to 100.0 mL with distilled water. Initial concentration was presumed to be in the vicinity of 10 molar, so this dilution was expected to give acid of somewhere around 0.1 molar.
Anhydrous sodium carbonate was prepared by heating the decahydrate at 150-160 °C for 3 hours in an oven. The dried material was stored in a vacuum desiccator until cool. The laboratory was at 48% relative humidity; mass change due to hydration was expected to be minimal during the few minutes the Na2CO3 was exposed to the air.
A 0.10 molar sodium carbonate solution was prepared by dissolving 2.650 grams of anhydrous Na2CO3 (fwt. 105.99) in enough distilled H2O to make 250.0 mL of solution.
The indicator chosen was methyl orange, which happens to have its color midpoint at the second equivalence point of the HCl / sodium carbonate titration curve. This color change is at pH 3.7 and goes from yellow to orange.
| Calculations: pH at the equivalence point Methyl orange has its midpoint at pH 3.7. The sodium carbonate-hydrochloric acid titration curve has two equivalence points. The first is where there is 1 mole Na2CO3 per 1 mole HCl; at this stage the carbonate would be virtually all converted to bicarbonate. The second equivalence point is where there is 1 mole Na2CO3 for every 2 moles of HCl; the main species here is H2CO3. The pH therefore depends primarily on how much H2CO3 ionizes. We started with a solution that was 0.1 M in Na2CO3, but at the second equivalence point it is effectively 0.1 M in H2CO3. To calculate the pH, the question becomes "what is the pH of a 0.1 M solution of H2CO3?" H2CO3 <------> H+ + HCO3- Ka = 4.3 x 10-7 = ([H+][HCO3-]) / [H2CO3] [H+] = [HCO3-] = x [H2CO3] = 0.1 - x 4.3 x 10-7 = x2 / (0.1 - x) x2 + (4.3 x 10-7)x - (4.3 x 10-8) = 0 x = 2.07 x 10-4 = [H+] pH = 3.68 This happens to be just about the midpoint for methyl orange indicator. |
5.00 mL (+/- 0.015 mL) of the 0.100 M sodium carbonate solution were put in a small beaker using a volumetric pipette. To this was added a carefully-measured amount of the unknown (ca. 0.10 M) HCl solution until a color change was noted. A high precision micro burette (capacity 5.00 mL) was used to deliver the solution; three fillings were necessary to provide the necessary amount of HCl, the equivalence point being reached about halfway through the third aliquot of 5 mL. We kept track of the exact amount required to reach color-change.
Localized color change began to appear around 7.40 mL of added HCl, becoming pronounced by 12.40 mL; this went away with swirling. The complete change-over of the solution from orange to red-orange happened at 13.73 mL.
Since there were calculated to be 5.00 x 10-4 moles sodium carbonate in the beaker, and since the methyl orange equivalence point was deemed to occur where there were 2 moles HCl per 1 mole Na2CO3, there must therefore have been 10.00 x 10-4 moles HCl added to the beaker.
0.00100 moles HCl in 13.73 mL comes out to 0.0728 moles per liter. Since the original HCl sample was diluted 1 to 100, the concentration of HCl in the original sample was 7.28 moles / liter.
In another trial, the titration was found to require 14.14 mL of HCl, for a concentration of 7.07 moles / liter.
Discussion:
The micro-apparatus for purifying HCl was found to give 7.1 to 7.3 M HCl.
The titration methods used in this experiment led to at least a 3% margin
of error; at least some of this can be attributed to the failure of
methyl orange indicator to have a sharp color-change. Time permitting,
the ideal way to do this experiment would be (1) to use methyl orange + xylene
cyanol (OR use a different base standard and indicator), and (2) to conduct
multiple trials to assess the relative deviation and the mean.
The molarity of HCl determined in the present study reflects the concentration
obtained from three, combined distillation runs. Since 6 M is strong
HCl and is a standard laboratory concentration, the product of the micro still
is more than strong enough for general lab use.
When 37% hydrochloric acid is heated, HCl gas escapes until the solution
reaches about 20.2%. This azeotropic solution is about 6 M in HCl and will
distill over in this same concentration. The final product of the micro
still is therefore a somewhat enriched version of this.
It is possible that procedural modifications can enrich the acid beyond
the 7.1-7.3 M that was obtained in this experiment. Since no vapor
escapes the still until the seal is unwrapped, it is possible that concentration
can be increased by:(1) redesigning the apparatus to have less internal air space in order to reduce the amount of undissolved HCl gas,
(2) refluxing the HCl from the micro beaker onto the cold funnel for a longer time, making sure to keep fresh ice in the funnel, and
(3) cooling the entire, sealed apparatus in an ice bath for perhaps 24 hours to help HCl gas pass back into solution before opening the system to the outside air.
If no excess HCl remains outside the micro beaker, the concentration of
HCl inside the micro beaker can theoretically approach 12 M, lessened by
whatever amount of HCl gas remains undissolved or tied up in any droplets
that remain on the inside wall of the beaker. Since the cold funnel
sends liquid on a one-way trip into the micro beaker, prolonged operation
of the system should be able to get fairly close to 12 M final concentration.
10 M HCl at least should be within reach.
If the micro beaker's maximum capacity is 15 mL, there should be no more
than 15 mL of impure HCl added to the system at the outset. This is
important to ensure maximal concentration of the acid.
Methyl orange is a useful indicator for the sodium carbonate-HCl titration,
but the equivalence point is not very sharp. Mixing it with some xylene
cyanole FF indicator, as is sometimes done in analytical labs, provides a
narrower color transition, although the colors involved are not the same.
Because Ka values that appear in books were derived experimentally
at some point, it is impossible to calculate the equivalence-point pH of
a titration with 100% confidence. The Ka values in different
sources may vary by as much as ~20%, even for the same listed concentration
and approximate temperature. If we assume +/- 20% error in the calculated
hydrogen ion concentration, that translates to an uncertainty of about 1/10
of a pH unit.
In the present experiment, the value 4.3 x 10-7, obtained from
the CRC Handbook of Chemistry and Physics (1988), was used.
Anhydrous Na2CO3 is often designated and sold as a
primary base standard; however, this is a puzzling choice when one considers
that anhydrous Na2CO3 absorbs moisture from the air.
Hygroscopic compounds do not make good primary standards; their mass
and composition can change during a laboratory session. Furthermore,
using carbonate in a titration adds an extra layer of complexity due to buffering.
A much better primary base standard is barium hydroxide. Nevertheless,
passable results can be obtained with sodium carbonate if the lab is kept
at low humidity and the compound is handled in such a way as to keep air exposure
to a minimum.
This experiment has shown a simple, convenient method for purifying small
amounts of hydrochloric acid. Future studies may refine the method
to enhance the concentration of HCl. While no trace metal analysis was
performed on the distillate, it is safe to assume the vast majority of metal
ions are removed during sub-boiling distillation.References
Pavia, D., Lampan, G., and Kriz, G. Introduction to Organic Laboratory Techniques. Philadelphia: Saunders College Publishing, 1988.
Wagner, R. A Handbook of Chemical Technology. New York: Appleton & Co., 1872.
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