Investigation of the Carbonate /Bicarbonate system.
Letort Vanessa
Student number: 2029474
HEV 4031
Kershaw Paul
ALKALINITY OF AQUEOUS SYSTEMS
INVESTIGATION OF THE CARBONATE /BICARBONATE SYSTEM
AIMS
This laboratory session should enable the operator:
. To define and understand the importance of alkalinity of aqueous systems.
2. To determine what constitutes a buffer system.
3. To understand the buffering capacity of chemical species such as carbonate and bicarbonate.
4. To gain knowledge of the carbonate/bicarbonate system and its environmental importance in water bodies.
5. To determine carbonate and bicarbonate equivalence points.
6. To practice two different titrations, pH and colour indicators titrations.
7. To calculate the concentration of carbonate and bicarbonate in an aqueous system.
8. To compare accuracy of both titrations.
9. To investigate, identify, and assess error limits associated with those procedures.
INTRODUCTION
Carbon is present in natural waters as the carbonate ion, the bicarbonate ion, and as dissolved carbon dioxide.
Carbonate ions come from the weathering of carbonate rocks, leaching in surface waters.
Carbon dioxide is present as a gas in the atmosphere. Carbon dioxide is a product of respiration of plants and animals and is released into our atmosphere when carbon-containing fossil fuels are burned in air. It makes up for only 0.03% of the total mass of the atmosphere. This fraction seems small but the function of CO2 in global biological and environmental systems is of prime importance (Loewenthal & Marais, 1978). Despite its low atmospheric concentration, carbon dioxide is normally abundant in natural waters because its solubility is more than 30 times higher than that of oxygen. The marine carbonate system represents the largest carbon pool in the atmosphere, biosphere, and ocean and is therefore of primary importance for the partition of excess carbon dioxide produced by man.
Carbon dioxide in the water dissolves, making carbonic acid, which lowers the pH. The pH can rise due to removal of carbonic acid by plants and algae. In water, the pH is mainly controlled by the concentration of CO2 present. Free hydrogen ions lower the pH, increasing the acidity of the water. The pH directly affects aquatic organisms. Extreme highs and lows in pH can cause damage to gill tissues. Respiration causes increased CO2 concentration and the pH to drop. The table below (table 1) gives some special effects of pH on fish and aquatic life. The Kentucky River Basin as part of their assessment report publishes these informations (Internet 4).
Limiting pH Values
Minimum
Maximum
Effects
3.8
0.0
Fish eggs could hatch, but deformed young were often produced
4.0
0.1
Limits for the most resistant fish species
4.1
9.5
Range tolerated by trout
4.3
---
Carp died in five days
4.5
9.0
Trout eggs and larvae develop normally
4.6
9.5
Limits for perch
5.0
---
Limits for stickleback fish
5.0
9.0
Tolerable range for most fish
---
8.7
Upper limit for good fishing waters
5.4
1.4
Fish avoided waters beyond these limits
6.0
7.2
Optimum (best) range for fish eggs
.0
---
Mosquito larvae were destroyed at this pH value
3.3
4.7
Mosquito larvae lived within this range
7.5
8.4
Best range for the growth of algae
Table 1: Effects of pH on aquatic life.
The main effect of pH is on the solubility and bioavaibility of run off substances in water (iron, lead, chromium, ammonia, mercury and so on). Those substances come from diverse anthropogenic activities (agricultural, domestic, industrial).
As the pH falls, (solution becomes more acidic) many insoluble substances become more soluble and thus available for absorption. For example, 4 mg/L of iron would not present a toxic effect at a pH of 4.8. However, as little as 0.9 mg/L of iron at a pH of 5.5 can cause fish to die. Since pH has a direct effect on organisms as well as an indirect effect on the toxicity of certain other pollutants in the water, maintaining a relatively constant pH is important to aquatic life.
The pH of natural water is in a dynamic equilibrium. There are many factors exerting an influence on the pH, and these are counteracted by the buffer system (Porteous, 2000). In natural waters, the carbonate/bicarbonate system is the most important buffer system consisting of a carbon dioxide, water, carbonic acid, bicarbonate, and carbonate ion equilibrium. The buffering capacity of an aqueous system can be defined as the ability of the water to resist changes in pH. It is an important factor in the control of the uptake capacity of the oceanic reservoir. The buffering capacity is important to water quality.
The buffering ability of the carbonate/bicarbonate system in aqueous solution is best explained by the chemistry of the system. First, CO2 dissolves in water:
(1)
Then, an equilibrium is established between dissolved carbon and carbonic acid, it is the stage of dissolution of the ,
(2)
Carbon dioxide reacts with water to form carbonic acid (pH is less than 4.5), only about 1% of the dissolved CO2 exits as .
Then the weak acid carbonic acid dissociates in two steps:
First, carbonic acid donates one proton to create bicarbonate, HCO3- ions (pH between 4.5-8.3).
(3)
Then, bicarbonate donates the last proton to form carbonate, CO32- (pH is greater than 8.3).
(4)
At a pH below 4.5 no bicarbonate or carbonate are present. Below pH 8.3, no carbonate ions are present. Finally, above pH 8.3 carbon dioxide or carbonic acid are not present (Internet 2).
This system is a diprotic acid system. There are three soluble components. The first one is the carbonic Acid, H2CO3, which can donate two protons (a weak acid), the second is the bicarbonate, HCO3-, is amphoteric (can donate or accept one proton, acid or base), and the last one is the carbonate, CO32-, which can accept two protons (a base). Table 2 displays the three major chemical reactions associated with the carbonate system as well as the corresponding equilibrium constants (Internet 1).
Reaction
Explanation
Equilibrium constant (25oC)
CO2,g +H2O <====> H2CO3*
Dissolution of CO2
pKH = 1.5 (at 25oC)
H2CO3* <====> HCO3- + H+
First deprotonation stage
PK1 = 6.3 (at 25oC)
HCO3- <====> CO32- + H+
Second deprotonation stage
PK2 = 10.3 (at 25oC)
Table 2: Chemical Reactions in the Carbonate System (Internet 1).
Reactions (2), (3), and (4) can be described by Graph 1. This graph compares the three carbonate species through their mole fractions to one another at a certain pH level.
Figure 1: The effect of pH on the distribution of carbonate species in solution (Windfield, 1995).
At a lower pH, is the dominant species in the solution. Between pH=6.3 and pH=10.3, is the dominant species in solution, and above a pH=10.3, is the dominant species. HCO3- and CO32- are major sources of pH changes in solution.
Alkalinity is a measure of the buffering capacity of the water, a measure of the capacity of water to neutralize or "buffer" acids.
The carbonate system is a major producer of alkalinity in water systems since CO2 is quite abundant in the atmosphere of Earth. Bicarbonate, HCO3- can act as an acid or a base as conditions dictate (equations (3) and (4)). One carbonate ion will neutralize two hydrogen ions. Because of their ability to either accept or give up hydrogen ions, buffering ions provide resistance against sudden changes of pH in the water. However, this buffering capacity is not indefinite.
Alkalinity is the concentration of titratable bases in water and measured in mg/l of equivalent calcium carbonate. Therefore, alkalinity is the sum of the titratable bases (ions with the ability to neutralize a hydrogen ion). The three chemical species responsible for alkalinity are:
- bicarbonate, forming the bicarbonate alkalinity = [HCO3-],
- carbonate, forming the carbonate alkalinity = [CO32-],
- hydroxide, forming the hydroxide alkalinity = [OH-] (also called phenolphthalein alkalinity)
The total alkalinity is the sum of each alkalinity parameters, Total alkalinity= [HCO3-] + [CO32-] + [OH-].
In natural waters, the concentration ...
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Alkalinity is the concentration of titratable bases in water and measured in mg/l of equivalent calcium carbonate. Therefore, alkalinity is the sum of the titratable bases (ions with the ability to neutralize a hydrogen ion). The three chemical species responsible for alkalinity are:
- bicarbonate, forming the bicarbonate alkalinity = [HCO3-],
- carbonate, forming the carbonate alkalinity = [CO32-],
- hydroxide, forming the hydroxide alkalinity = [OH-] (also called phenolphthalein alkalinity)
The total alkalinity is the sum of each alkalinity parameters, Total alkalinity= [HCO3-] + [CO32-] + [OH-].
In natural waters, the concentration of hydroxyl and hydrogen ions are very low because water very weakly dissociates. Primarily, carbonate and bicarbonate are responsible for all the measurable alkalinity. Seawater has an average alkalinity of 116mg/l. Actually, there is no published alkalinity requirement (Internet 2).
Most oceanic waters have a pH in the range 8 to 8.3, as they contain more hydroxide ion than hydrogen ion due to:
(5)
(6)
Natural water unpolluted is alkaline (P. O'Neill, 1993)
Alkalinity levels will vary with geographical locations. Most alkalinity in surface water comes from calcium carbonate, CaCO3, being leached from rocks and soil. This process is enhanced if the rocks and soil have been broken up for any reason, such as mining or urban development (Porteous, 2000). Limestone contains especially high levels of calcium carbonate.
Some sample data are shown in Figure 2. In general, water in the eastern half of the United States will have a higher alkalinity than water in the west because of a higher occurrence of limestone. Areas in the extreme northeast that have had the limestone scoured away by glacial action will often have a lower alkalinity (Source: Internet 6).
Alkalinity is reported in units of mg/L CaCO3, because the carbonate ion, CO32-, is its primary constituent.
Site
Alkalinity
Mg/l CaCO3
Missouri River, St. Joseph, MO
224
Missouri River, Garrison Dam, ND
78
Cataloochee Creek, Cataloochee, NC
626
Columbia River, Northport, WA
49
Merrimack River, Lowell, MA
7
Figure 2: Alkalinity of selected rivers US.
Calculating the buffering capacity of an aqueous solution is not easy because natural waters are complex systems. However, determining the alkalinity of an aqueous system is relatively simple. Alkalinity is determined by titration of the water or the aqueous system studied with a strong acid (if the solution studied is a base and inversely for an acidic solution). The volume of acid used to reach the equivalence points, is a measure of the alkalinity of the solution. These equivalence points represent the equilibrium of equations (3) and (4).
Before the time when pHmeters came into general use, the equivalence points were determined approximately by using colour pH indicator such as methyl orange and phenolphthalein.
EXPERIMENTAL PROCEDURES
The experimental procedures involved during this laboratory session are fully explained in the handbook.
The whole class has been divided into three groups of students (two groups of two students, and one group of three students).
The pHmeter has been calibrated with a buffer solution at pH=10 instead of pH=9.2 as indicated in the handbook. The electrodes have been rinsed with distilled water between each reading.
Group 1 has undertaken a thorough cleaning of the glassware used, before experimental handling. Most of the glassware has been rinse with the carbonate/bicarbonate mixture to be titrated before titrations and pH readings (thus explaining the disappearance of most of the mixture during the experiment handling).
The titration with colour indicators has been undertaken before the pHmeter procedure in order to anticipate the pH drop for the second procedure.
PHENOLPHTALEIN AND METHYL ORANGE END POINTS
Methyl orange is an azo dye that changes colour from yellow to red as the pH is lowered below about 4.5. It is also known as p-dimethylaminoazobenzene-p'-sulfanilic acid.
Phenolphthalein is a polyphenolic compound, which loses both a water molecule and a hydrogen ion at high pH. As the protonation occurs, the colour changes from pink to colourless.
Phenolphthalein is commonly used indicator for titrations, and is a weak acid.
In this case, the weak acid is colourless and its ion is bright pink. Adding hydroxide ions removes the hydrogen ions from the equilibrium, which tips to the right to replace them - turning the indicator pink (Internet 3).
Phenolphthalein is a colour indicator that changes from pink to colourless at pH 8.3 when acid is added.
When determining carbonate and bicarbonate concentrations in an aqueous solution by titration, by convention, the pH is usually about 4.5 for the alkalinity titration and 8.3 for acidity. When the solution titrated reached pH=8.3, the concentration of the volume of acid added equal to the sum of concentrations of the base solution, the carbon dioxide and the bicarbonate originally present. This pH 8.3 is as well the approximate pH for which a colour change occurs for phenolphthalein, so that the volume of titrant required to reach that point is called the phenolphthalein alkalinity. pH 8.3 is the phenolphthalein endpoint.
For a sample initially containing only carbonate, such as the solution titrated during this experiment, the volume required to achieve the phenolphthalein colour change is close to half that required for the subsequent methyl orange end point. It is not exactly half because the colour change and the equivalence point are not exactly the same (Internet 3).
RESULTS
. Class results of the two procedures
Table 3 shows the volume of HCl added to the carbonate/bicarbonate solution, to obtain either a colour change with the titration with colour indicator, or to obtain the deepest change in pH.
VOLUME OF HCl ADDED TO THE SOLUTION (ml)
GROUP
INDICATORS TITRATIONS
PHmeter TITRATION
PHENOL
PHTALEIN
METHYL ORANGE
First deepest change
Second deepest change
GROUP 1
3.65*
3.65,3.7,3.62
2.8
2.9
3.4 to 3.6
averaged to 3.5
2.9
GROUP 2
4.15*
4.5,4.2,4.1
2.8
2.7
N/A
N/A
GROUP 3
3.8*
3.8,3.7,3.9,3.7
3.3
3.2
N/A
N/A
Table 3: Volume of HCl added (ml) during each procedure.
Results obtained during the pHmeter titration have not been shared, so that for this report only results from group 1 appear. The volume of HCl added during the first procedure varies between 3.65 ml and 4.15 ml for the first end point, and 12.7 ml and 13.3 ml for the second end point. The volumes of acid added for which the deepest changes in pH occur for the second procedure, have been determined by reading the derivative cure (Graph 2).
2. Alkalinity curve obtained using the pHmeter procedure (results of group 1)
Table 4 in Appendix displays all the readings that were obtained during the pHmeter procedure, after adding a known volume of HCl.
Graph 1 has been obtained by plotting the pH change resulting after adding different volume of HCl. The volume of HCl added is plotted on the x-axis and the resulting pH reading on the y-axis.
Graph 1: Diagram of the alkalinity curve obtained when pH changes are plotted against the volume of acid added to the 25 ml of carbonate/bicarbonate solution titrated.
Graph 1 shows the effects of all the carbonate species in the aqueous system. Five stages are present (AB, BC, CD, DE and EF).
The actual volume of HCl added, corresponding to end points of phenolphthalein and methyl orange, will be best shown by plotting a derivative curve for the data of Table 4, the table obtained is Table 5 in Appendix, which relates the variation in pH (A) and the variation in volume of acid added (B) during the procedure. Those latest results have been plotted (A/B on the y-axis and total volume of acid added on the x-axis) on Graph 2.
Graph 2: Derivative curve for accurate readings of the equivalence points of carbonate and bicarbonate
Using graph 2, two maxima values (X and Y) of volume of HCl added have been identified, and added to Table 3.
3. Calculation of carbonate and bicarbonate levels
* Carbonate
From the equation:
(7)
One mole of carbonate reacts with two moles of HCl. Knowing that the concentration of HCl is 0.1 M/ml i.e. 1*10 mole of HCl=1/2*10mole of =1/2*10*60=3 mg of for every unit of ml of HCl used. The sample used was of 25 ml; the results need to be multiplied by 40 to be expressed per litre.
* Bicarbonate
From the equation:
(8)
One mole of bicarbonate reacts with one mole of acid, knowing that the concentration of HCl is 0.1M / ml i.e. 1* 10 of HCl=1*10 mole of =1*10*61=6.1 mg of , for every ml of HCl used. Twenty-five ml of sample was used, so the results need to be multiplied by 40 to be expressed per litre.
The solution titrated reached a colour change with phenolphthalein as a colour indicator when half the carbonate has reacted with the acid to become bicarbonate; the volume of acid added to obtain this is noted x. Thus, twice as much carbonate is actually present in the solution and the volume obtained needs to be multiplied by two to obtain the level of carbonate present in the solution, this volume is 2x.
The final volume of acid added to obtain the methyl orange or the second steepest change in pH, is noted y. Thus, the volume is y-2x.
Using those results, calculations of the level of carbonate and bicarbonate can be undertaken.
Below, an example of those calculations is provided; every result has been calculated this way.
Example:
Group 2, referring to Table 3, have found an average of 4.15 for x, and an average of 12.75 for y.
Bicarbonate level: y-2x=12.75-2(4.15)=4.45 ml
ml of 0.1 M of HCl=6.1 mg of bicarbonate
Thus, for 4.45 ml of HCL, there are (4.45*6.1)=27.145 mg of bicarbonate.
For 1 litre, there are (22.145*40)=1.0858 g/l of , in the solution.
Carbonate level: 2x=2*4.15=8.3 ml
ml of 0.1 M of HCL=3 mg of carbonate
Thus, for 8.3 ml of HCl, there are (8.3*3)=24.9 mg of carbonate.
For 1 litre, there are (24.9*40)=0.996 g/l of , in the solution.
Table 5 displays the final results of level of carbonate and bicarbonate obtained after calculation, for both procedures, using Table 3.
GROUP
INDICATORS PROCEDURE
PHmeter PROCEDURE
Results for group 1 only
Carbonate
g/l
Bicarbonate g/l
Carbonate
g/l
Bicarbonate g/l
Group 1
0.876
.354
Group 2
0.996
.086
0.840
.440
Group 3
0.912
.379
Table 5: Level of carbonate and bicarbonate present in aqueous solution titrated, according to two different titrations.
Both procedures seem to indicate that the level of carbonate is lower than the level of bicarbonate. However, results between groups and between procedures differ.
DISCUSSION
. Equations, Graph 1 interpretations, carbonate and bicarbonate concentrations.
At the beginning of the titration, CO32- exists to the practical exclusion of the other carbonate species. When one equivalent of acid has been added, almost all of the CO32- has been changed into HCO3-. Addition of a further equivalent of acid changes practically all hydrogen carbonate into carbonic acid, H2CO3. The latter is in equilibrium with water and CO2.
In this laboratory experiment, results show (Table 5) that the concentration of carbonate, , is ranging between 0.876 and 0.996 g/l; and the concentration of bicarbonate, , is ranging between 1.0858 and 1.3786 g/l.
These levels represent the concentrations of these two species at the equilibrium points of equations (3) and (4).
Graph 1 shows that the carbonate/bicarbonate mixture investigated changes from a pH 9.8 to pH 2.1 with the addition of 20 ml of HCl (strong acid). During stage (AB), a drop of 2 units of pH is generated by addition of 3.4 ml of acid, whereas the same drop in pH unit occurs for an addition of 1 ml of HCl during stage (BC). During stage (CD), a drop of 1 unit in pH is generated by addition of 4 ml of acid, whereas the same drop in pH is observed for an addition of 0.8 ml of HCL.
(AB) and (CD) are stages where the solution is able to withstand relatively large additions of acid without large variation of pH.
Stage (CD) in Graph 1 represents the ability of the system to neutralize acid; the longer it is the more stable the chemistry of the water is. This means that for a long (CD) stage more acid needs to be added to reach (DE). And most importantly, no significant pH changes occur for the flora and fauna. During (CD), 4 ml of acid is needed to obtain 1 unit change in pH. This ability of a system to neutralize the effect of acid or base added, is its buffering capacity.
Peter O'Neill, from the department of environmental sciences at the University of Plymouth, stated in 1993 that stages (BC) and (DE) occur in natural water systems at pH 8 to 8.5 and pH 4 to 4.5, respectively. This statement confirms the results found during the experiment as Graph 1 shows that the range of pH of (BC) is between 9 and 6.8, and (DE) is ranging between pH 5 and 2.6. During stage (DE), 0.7 ml of HCl is added to obtain a change of one unit in pH.
However, the solution titrated was not a natural water sample but a prepared solution. Thus, both stages are not included in such a narrow range of pH.
Table 5, finally, displays the concentrations of carbonate and bicarbonate. Those range between 0.840 g/l and 0.996 g/l for the carbonate, and between 1.086 g/l and 1.440 g/l for the bicarbonate. For protection of aquatic life, the buffering capacity should be at least 20 mg/L. If alkalinity is naturally low, (less than 20 mg/L) there can be no greater than a 25% reduction in alkalinity (Winfield, 1995).
2. Comparison of the two procedures
COLOUR INDICATORS TITRATION
PHmeter TITRATION
* Procedure fast to execute
* Possibility of repeating the titration in the time allowed, so that results can be averaged
* Small amount of calculations need to be done
* Difficulty to determine the volume of HCl for which a colour change occurs
* Colour change not obvious
* Procedure requesting small amount of material (glassware), thus less error due to equipment
* No calibration of equipment to be done
* Two different indicators need to be used for each end points
* Procedure relatively inaccurate
* Procedure is time consuming
* Lots of readings need to be undertaken to plot a curve
* Lots of calculations to be done
* Procedure could be undertaken only once in the time allowed
* Only one set of results obtained, thus, no comparison of results within the procedure
* Change of pH immediately visible on the pHmeter
* Procedure requesting the use of lots of glassware, thus, more error could be due to equipment
* The pHmeter needs to be calibrated
* Only the pHmeter is used to determine the end points
* Plotting a graph and a derivative curve is necessary, most of the results depend on calculations
* Procedure relatively accurate
Table 6: Comparison of the two procedures.
Both methods are acceptable, as far as the accuracy of results is concerned. Table 5 indicates that whatever the procedure, the carbonate concentration is higher than the bicarbonate concentration.
However, the colour indicators method was routinely used before the pHmeter procedure was developed. Then, one had the choice of procedures. Depending on conditions of experiment, one or the other method is better to be used. For example, the potentiometric method (pHmeter) can be difficult to employ whilst working in the field, as an electrical power is needed to use the pHmeter. Nevertheless, it is generally accepted that the potentiometric method is more accurate than the colour indicators method (Golterman, Clymo & Ohnstad, 1978)
As indicating in Table 6, the pHmeter procedure is more time consuming than the colour indicator procedure, plus it requires graph analysis skills. It has also been noticed that the pHmeter method relies on proper operation of the pHmeter, which is a sensitive instrument.
In addition, it has been observed (Barnes, Ivan. 1964) that interference (surfactants and precipitates which coat the electrode) can obstruct results.
In the other hand, the obvious disadvantage of the colour indicator procedure is that it relies mainly on observing the colour change, which is relatively subjective.
3. Comparison of results
a) Comparison between groups
The solution of carbonate titrated came from the same sample for each group for this part of the experiment.
Comparison of results obtained between groups, can only be done within the first procedure, as no results were shared for the potentiometric titration.
The variation range between groups within the indicator titration procedure is 0.88 ml (minimum vol. of acid added minus maximum vol. of acid added), difference between group 1 and group 2 being the greatest, to reach the phenolphthalein end point. The variation in volume to reach the methyl orange end point is 0.6ml. These variations could be insignificant, but to determine the relevance of the results variations, we need to consider the range of volume for which a change occurs and compare it to the differences found in results.
Looking at Graph 1, the range of volume for which a sudden change in pH occurs when adding between 3 ml and 5 ml of HCl and between 12 ml and 14 ml of HCl for the second end point. Thus, both ranges have a variation of 3 ml. The range of differences for the both end points represent 20% to 30% of the range of volume. These percentages are quite high, and it will be considered that errors have been done during handling of the experiment.
Compare to other groups, group 2 found that a higher volume of HCl needed to be added to obtain a phenolphthalein colour change, and group 3 found that a larger volume of HCl needed to be added to obtain a methyl orange colour change.
This reveals differences in way of undertaking the experiment between groups. However, it should be noticed that the number of results it self, does not permit a relevant investigation. The bigger the number of results the more accurate is the analysis (a minimum of 10 is necessary).
Many factors can affect the experiment thus results.
The first one is the number of times the procedure has been repeated. As seen on Table 3, all the groups have repeated the titration with phenolphthalein 3 times or more and twice for the titration with methyl orange. Thus, the variation in results between groups was not affected by the time allowed or by the number of results.
Another factor, and certainly the more relevant, is the ability of the operators to spot the colour change of the indicator, especially for the methyl orange indicator. One burette drop is approximately 0.05 ml of HCl, after calculations is seems that one burette drop of HCl can induce an error of 0.096% in the final result of bicarbonate concentration. This mean that to get a percentage of difference of 20%, Group 2 would have had to pour an excess of 1.2 ml (24 drops) of HCl compare to Group 1 and 3. This seems, a priori, unlikely to happen when undertaking an experiment. However, the colour change of the methyl orange is difficult to pin point, and one can argue when the solution change from a pale yellow to a pink coloration. Here, this factor will thus be considered as the most relevant one, and it will be advised to use a blank solution to compare colour change or to use another colour indicator for the titration of carbonate. Bromocresol green could be used. Bromocresol green however requires the operator to boil the solution and then to cool it down (making it very difficult to use for fast on site titration) to remove the CO2 making the procedure become the titration of a strong base with a strong acid, and consequently improving the accuracy (0.1%) of the method (Internet 5).
The phenolphthalein colour change is easier to detect but still inaccurate. Differences in results between groups should be considered due to variation of detection of the end point, thus a difference in value of acid added. It was also noticed by Skougstad & Fishman (1979) that titration with phenolphthalein (to pH 8) by acids and titration by weak acids will give incorrect results if much carbonate is present.
Finally, the rinsing the glassware used (pipettes, beakers) with distilled water or the solution it will support, would help to reduce effects of alien chemicals species. Although, this glassware is cleaned up between experimental sessions, so that any contaminant chemical species will be of very low concentrations. Considering that concentration levels of the species investigated are not trace amount, this factor is though to have a low impact on results,
b) Comparison of results between procedures
Only one set of results has been obtained during the potentiometric procedure.
The averaged volume of HCl added will be used for the comparison. Each group has executed the first titration at least twice. Table 3 displays averaged volume of HCl added, as well as individual values, for the phenolphthalein end point.
According to Table 3, volumes of acid used to reach both end points and the steepest changes in pH, are different.
The averaged volume of HCl added to obtain the end point of phenolphthalein is 3.86 ml for the first procedure, and 3.5 ml (averaged) of HCl have been titrated to reach the first end point with the second method. The difference in results is thus, 0.36ml.
The averaged volume of HCl added to obtain the second end point is 12.95 ml, and 12.90 ml for the second procedure. Thus, the difference in results is 0.05 ml of HCl.
The significance of the differences is investigated below.
As previously done, a >10% difference will be considered as significant. The first end point volumes (i.e. 3.86 ml and 3.5 ml) differ of 9.78% and the second end points values (i.e. 12.95 ml and 12.9 ml), differ of 0.39%. It seems that titrating for bicarbonate with one or the other procedure, does not affect results (less than 1% difference), whereas when titrating for carbonate, the results will vary of around 10% between each other.
Errors could have occured during calculations inherent of both procedures, a mishandling of the instruments (calibration of the pHmeter, titration and so on), a lack of data to determine the end points of the reactions, as well as reading errors and contamination of glassware.
It has been noticed previously that the potentiometric procedure is accepted to be the most accurate method. The difference in results, here, could be due to this fact. As cited above (section 3.a)), using colour indicators results in inaccuracy of results mainly due to the personal characteristics (colour change spotting).
Moreover, the carbonate solution ran out during the experiment. The laboratory technician had to quickly produced another carbonate/bicarbonate mixture for the experiment to carry on. Thus, the mixture titrated during the pHmeter procedure was different from the solution titrated during the colour indicators procedure. Although, the second mixture prepared has the same species concentration than the first one, and was supposed to be chemically identical, it makes it quite difficult to identify errors due to handling of the experiment.
In the case of these two procedures, errors are significant to results, which means that the degree of errors done during the titrations, the calculations or the graphic reading, is important enough to change the final results.
c) Factors affecting procedures
The potentiometric method requires the use of a magnetic stirrer, which if selected at maximum speed, might allow further gaseous exchanges with the ambient atmosphere, allowing more carbon dioxide to dissolve. Loss of carbon dioxide can also occur during titration, and quantifying this loss is difficult (Faust & Aly, 1981). Carbon dioxide exchange happens if the partial pressure of this molecule between water and air phases, is different.
However, loss or gain of carbon dioxide can be considered as negligible providing a smooth stirring of the solution (Weber and Stumm, 1962).
Loewanthal and Marais (1978) also remarked that the selection of equivalence points is important. The carbonate equivalence point is affected by temperature, ionic strength, and the total species concentration. The bicarbonate equivalence point is affected by both temperature and ionic strength. Ionic strength (µ or I) is a measure of the total concentration of ions in a solution. Temperature affects the chemical species' dissociations constant of the carbonate/bicarbonate system (Table 7).
Reaction
H2CO3 <====> HCO3- + H+
PK1
HCO3- <====> CO32- + H+
PK2
Equilibrium constant at 25oC
PK1 = 6.37
PK2 = 10.33
Equilibrium constant at 90?C
PK1 = 6.33
PK2 =10.13
Table 7: Effect of temperature on equilibrium constant of carbonate and bicarbonate (source: Loewanthal & Marais, 1978)
Thus, operators had to assume that the solution titrated was ideal for the results to be representative of the sample. In an ideal solution, each particle behaves independently of any other particle, and thus hold only for very dilute ionic solutions.
CONCLUSION
During this session of experiment, students have appreciated two different ways of undertaking an experimental work.
It has been shown that:
. Natural waters possess the ability to neutralize acid, it is the alkalinity.
2. Alkalinity is defined by the concentrations of carbonate and bicarbonate of an aqueous system.
3. These concentrations are important to the stability of the water.
4. The alkalinity of an aqueous system can be investigated through two procedures, the colour indicator titration and the pHmeter titration.
5. During both titrations, the equilibrium points of equations (3) and (4) are determined, and through calculation and graphic reading the concentrations of both carbonate species can be determined.
6. The potentiometric procedure has been identified as being more accurate in results than the colour indicator titration.
7. Both procedures revealed their advantages and disadvantages.
8. Errors in laboratory are mainly due to the difficulty of identifying the colour change of the indicators, and the lack of data to plot a relevant alkalinity curve.
APPENDIX
Table 4: Readings of pH when adding HCl to the solution titrated. Method 2.
Vol. of HCl added
ml
PH reading
Vol. of HCl added
Ml
PH reading
0
9.8
0
5.6
9.5
1
5.4
2
9.2
2
5
3
8.6
2.2
4.8
3.1
8.45
2.3
4.7
3.2
8.35
2.4
4.65
3.3
8.2
2.5
4.55
3.4
8
2.6
4.4
3.5
7.8
2.7
4.2
3.6
7.6
2.8
4
3.7
7.5
2.9
3.6
3.8
7.4
3.
3.45
3.9
7.4
3.1
3.2
4
7.25
3.5
2.8
5
6.8
4
2.6
6
6.4
5
2.4
6.5
6.3
6
2.3
7
6.3
7
2.2
8
6.1
8
2.2
9
5.8
20
2.1
Table 5: Change of pH and change of volumes to plot derivative curve
Initial pH
Change in pH: A
Change in volume: B
A/B
Total volume of HCl used ml
Initial pH
Change in pH: A
Change in volume: B
A/B
Total volume of HCl used ml
9.8
0
0
0
0
5.6
0.2
0.2
0
9.5
0.3
0.3
5.4
0.2
0.2
1
9.2
0.3
0.3
2
5
0.4
0.4
2
8.6
0.6
0.6
3
4.8
0.2
0.2
2.2
8.45
0.15
0.1
.5
3.1
4.7
0.1
0.1
2.3
8.35
0.1
0.1
3.2
4.65
0.05
0.1
0.5
2.4
8.2
0.15
0.1
.5
3.3
4.55
0.1
0.1
2.5
8
0.2
0.1
2
3.4
4.4
0.15
0.1
.5
2.6
7.8
0.2
0.1
2
3.5
4.2
0.2
0.1
2
2.7
7.6
0.2
0.1
2
3.6
4
0.2
0.1
2
2.8
7.5
0.1
0.1
3.7
3.6
0.4
0.1
4
2.9
7.4
0.1
0.1
3.8
3.45
0.15
0.1
.5
3.
7.4
0
0.1
0
3.9
3.2
0.25
0.1
2.5
3.1
7.25
0.15
0.1
.5
4
2.8
0.4
0.4
3.5
6.8
0.45
0.45
5
2.6
0.2
0.5
0.4
4
6.4
0.4
0.4
6
2.4
0.2
0.2
5
6.3
0.1
0.5
0.2
6.5
2.3
0.1
0.1
6
6.3
0
0.5
0
7
2.2
0.1
0.1
7
6.1
0.2
0.2
8
2.2
0
0
8
5.8
0.3
0.3
9
2.1
0.1
0.1
20
REFERENCES
> Faust, D. S. and Aly, M. O. 1981. Chemistry of natural waters. Ann Arbor Science Publishers Inc.
> Golterman, H. L., Clymo, R. S., and Ohnstad, M. A. M. 1978. Methods for Physical and Chemical Analysis of Fresh Waters. Second Edition. Blackwell Scientific Publications.
> Porteous, A. 2000. Dictionary of Environmental Science and Technology. (Third Editon). John Wiley & Sons, Ltd.
> Winfield, A. 1995. Environmental Chemistry. Cambridge University Press.
> Loewanthal, R. E., and Marais, G. R. 1978. Carbonate Chemistry of Aquatic systems. Theory and Application. Second Edition. Ann Arvor Science Publishers Inc.
> Weber, W. J. and Stumm, W. 1963. Mechanisms of Hydrogen Ion Buffering in Natural Waters. J. A. W. W. A.
> Skougstad, M., Fishman M., Friedman, L.C., Erdman, D.E. and Durran, S.S. (editors) 1979. Methods for determination of inorganic substances in water and fluvial sediments. Techniques of Water-Resources Investigations of the United States Geological Survey, book 5, chapter Al, p. 626
> Barnes, Ivan. 1964. Field measurement of alkalinity and pH. U.S. Geological Survey Water-Supply Paper 1535-H, p. 17.
Internet references:
.
> Internet 1: The Carbonate System
http://www.swbic.org/education/env.engr/carbonate/carbonate.html
> Internet 2: Water Quality
http://www.hatcherybiotech.com/water.html
> Internet 3: Chapter 16 Volumetrics Methods
http://www.ecs.umass.edu/cee/reckhow/courses/5721572bk16/
> Internet 4: Water Quality parameters
http://www.uky.edu/WaterResources/Watershed/KEB_AR/
> Internet 5: Determination of Carbonate Content of a Soda ash Sample
http://www.csudh.edu/olivier/che230/labmanual/carbonate.html
> Internet 6: CBL: Water Quality Alkalinity
http://www.euphrtaes.wpunj.edu/faculty/partnership/CBL/ex-cbl05.html