Investigation of the Carbonate /Bicarbonate system.

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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.
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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 ...

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