Development Team

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Environmental Sciences

Environmental Chemistry

Carbon Dioxide and Carbonate System Paper No: 16 Environmental Chemistry

Module: 16 Carbon Dioxide and Carbonate System

Development Team

Principal Investigator

&

Co- Principal Investigator

Prof. R.K. Kohli

Prof. V.K. Garg &Prof.AshokDhawan Central University of Punjab, Bathinda

Paper Coordinator

Prof. K.S. Gupta

University of Rajasthan, Jaipur

Content Writer

Prof. K.S. Gupta

University of Rajasthan, Jaipur Content Reviewer Dr. V.K. Garg

Central University of Punjab, Bathinda

Anchor Institute Central University of Punjab

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Carbon Dioxide and Carbonate System

Description of Module

Subject Name Environmental Sciences Paper Name Environmental Chemistry Module Name/Title

Carbon Dioxide and Carbonate System Module Id EVS/EC-XVI/16

Pre-requisites A basic knowledge of Environmental Chemistry

Objectives

1. To know about sources of CO2

2. To know about trends in CO2 emission 3. To know about sinks of CO2

4. To know about dissolution of CO2 in water 5. To know about CO2 in equilibrium with ocean

6. To know about effect of increase in CO2 on dissolution in oceans 7. To know about carbonate solubility and alkalinity

8. To know about carbonate solubility in a system open to air

Keywords

Sources of CO2, CO2 emission, sinks of CO2, CO2 in equilibrium with ocean, effect of increase in CO2 on dissolution in oceans, carbonate solubility and alkalinity

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Module 16: Carbon Dioxide and Carbonate System

Introduction

Carbon dioxide is one of the most important trace atmospheric constituents and has plays an important role in nature of climate prevailing in Earth and other planets. It is responsible for maintaining the average temperature of ~288 K (15 ̊ C) of the Earth , which makes it habitable.

Another important role is in photosynthesis in plants, which provide oxygen to our atmosphere and food for themselves and the mankind.

Contents 1. Introduction 2. Sources of CO2

3. Land use, land use change, agriculture and forestry 4. Incineration of waste products

5. Trends in CO2 emission 6. Sinks of CO2

7. Dissolution of CO2 in Water 8. CO2 in equilibrium with ocean

9. Effect of increase in CO2 on dissolution in oceans 10. Carbonate solubility and alkalinity

11. Carbonate solubility in a system open to air 12. Alkalinity

13. Suggested Reading

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Sources of CO2

Both natural and anthropogenic, i. e., manmade sources are responsible for the release of carbon dioxide in the atmosphere.

Natural Sources

Earth’s oceans, plants, animals, volcanoes, respiration by living aerobic organisms all are source of CO2. If the nature has CO2-sources, it has sinks also for the removal of CO2 from the atmosphere. These sources and sinks maintain a balance in CO2 concentration. This balancing has been going on for past several thousand years or more. Following are natural sources of CO2.

Oceans: The huge amount of CO2 is dissolved in different types of waters existing on Earth surface..

The atmospheric CO2 is in equilibrium with dissolved CO2 in oceans (1) .

CO2 (gas) CO2 (ocean) (1)

The ocean-atmosphere exchange of CO2 is the largest natural source. It contributes about 42.8% CO2

to the total CO2 emitted by natural sources.

Respiration by Plants and Living Species: Plants and animals need energy and they produce it by respiration, during which biochemical reactions occur which release CO2. The process of respiration contributes sizably – about 28.5% of total natural CO2 emission.

Soil: The respiration by bacteria, fungi, plants and animals living in soil is called soil respiration. The organisms present in soil biodegrade the dead organic matter. Both these processes release significant CO2.

Anthropogenic Sources

Beginning with the industrial revolution, there has been a continuous rise in anthropogenic emission of CO2. Incidentally, industrial revolution began with the invention of steam engine. Most of CO2 sources relate to energy production for industrial/domestic use and for the manufacture of goods.

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Fossil Fuel Combustion

Carbon, which has been lying buried in the Earth for millions of years in the form of coal, oil, natural gas, shell gas, etc, is being mined and used for production of energy through its combustion in thermal power plants, automobiles, rail, road, marine, air, industries, metallurgical processes, cement manufacture, burning fuels for heating and cooking, etc. By far, combustion of carbon based fuels is the largest manmade source CO2 Intergovernment Panel on Climate Change(IPCC, 2007) estimated the fossil fuel combustion to contribute about 57% of total manmade emission of CO2.

Land Use, Land Use Change, Agriculture and Forestry

With increasing urbanization and increasing demand for food and homes, land use change is taking place. The land, which as a natural environment, was acting as a natural buffer for CO2 is being used for construction of houses, factories, industries, roads, etc. Trees are carbon banks, absorb CO² from the atmosphere and store it. Thus, removal of plants/forests has two effects. One, the carbon, which was stored, is released and it will soon come in the form of CO2 in the atmosphere. Two, the removal of trees removes the carbon removal process from the atmosphere also. These processes of deforestation and decay of biomass, etc., are estimated to account for 17% of total manmade CO2

emission. Agricultural emissions are mainly concerned with management of soils, biomass burning, etc.

Incineration of Waste Products

Incineration of waste products such as plastics, textiles, etc, is a minor source of CO2 emission.

Trends in CO2 Emission

By collecting ice samples at different depths, the concentration of CO2 has been estimated in past thousand years from now. About twenty thousand years ago CO2 was ~200 ppm. Since then it has been rising, and has reached a level 280 ppm at the end of 19^th century. The carbon dioxide is being measured continuously at Mauna Loa, Hawaii, since 1958. In the past one hundred years, a trend of rapid rise in CO2 is seen. The concentration reached 315 ppm in 2000 and 380 ppm in 2008. The current average CO2 is 409.01 ppm as on January, 2018.

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The rising trend in CO2 is seen in the Mauna Loa CO2 Data Figure 1 also. According to US EPA, the percent contribution of some countries to total CO2 is: China = 23%, USA = 19%, European Union (excluding Estonia, Latvia, Lithuania) = 13%, India = 6%, Japan = 4%, Canada = 2% and others = 28%.

Sinks of CO2

A carbon dioxide sink is a reservoir in which CO2 is stored indefinitely. The sink can be natural or artificial. Sink continues to accumulate CO2. The process of removal of CO2 from atmosphere is known as carbon sequestration. The most vital and largest natural sinks are oceans and the photosynthesis by plants. The oceans have the huge capacity of absorption of CO2 governed by the equilibrium and aided by physicochemical and biological processes. Since increase in CO2 is directly linked with global warming, effort are being made to develop artificial carbon dioxide sequesters, which will capture and store CO2. However, so far there is not much success in this endeavor.

Figure : Atmospheric CO2 concentration at Mauna Loa observatory Source: https://www.dieselnet.com/news/2016/03noaa.php

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Dissolution of CO2 in Water

(An understanding of the following material would be greatly helped by going through the Modules 2- 5)

The solubility of carbon dioxide in water is governed by Henry’s law, which states that at a constant temperature, the amount of the gas dissolved is proportional to pressure of the gas(or partial pressure if the gas under question, if it is in a mixture of gases, such as air). This law is very helpful in understanding the solubility of atmospheric CO2 in all kinds of natural waters such freshwater, seawater, lake water, river water, etc. The dissolution of CO2 in water undergoes several equilibria and the presentation differs from books to books.

Henry’s law

Henry‘s law constant for CO2 dissolution equilibrium is defined as:

KH

CO2(g) + H2O CO2.H2O(H2CO3(aq)) (2)

KH = [H2CO3(aq)] / PCO2 (3) In Eq. (3), KH = Henry’s law constant in units of mol L^-1 atm^-1, [H2CO3] = concentration of CO2 dissolved in water, and PCO2 = atmospheric partial pressure of CO2 in unit of atm (atmosphere) at a specified temperature. The value of KH is 3.38 × 10^-2 mol L^-1 atm^-1 at 25^oC.

The H2CO3 undergoes two consecutive dissociation equilibria as shown in Eqs. (4-5).

Ka1

H2CO3 H⁺ + HCO3- (4) Ka2

HCO3- H⁺ + CO32- (5)

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HCO3- species is bicarbonate or hydrogen carbonate ion. Species CO32- is carbonate ion. The value of Ka1 = 4.4×10⁻⁷.

The bicarbonate ion dissociates into the carbonate ion, CO32−: Ka2

HCO3− CO32− + H+ ; Ka2 = 4.69×10⁻¹¹ mol/L; pKa2 = 10.329 (6)

The total concentration of CO2 dissolved in water, [CO2(aq)]Total , is given by Eq. (7).

[CO2(aq)]Total = [H2CO3] + [ HCO3−

]+ [CO32−

] (7)

The values of [H2CO3], [ HCO3− ] and [CO32− ] in terms of Eqs. (4 - 6) are as follows.

[H2CO3] = KHPco2 (8) [HCO3−] = Ka1[H2CO3] [H⁺] (7)

[CO32−] = Ka1Ka2KHPCO2/[H⁺]2 (9)

Substituting the values of [H2CO3], [HCO3−], [CO32− ] in Eq. (7), we get:

[CO2(aq)]Total = KHPCO2 + Ka1KHPCO2/[H⁺] + Ka1Ka2KHPCO2/[H⁺]2 (10) Equation(10) can be written as Eq. (11) or (12)

[CO2(aq)]Total = KHPCO2 [1 + Ka1/[H⁺] + Ka1Ka2/[H⁺]²] (11) = KHPCO2[ [H⁺]² + Ka1[H⁺] + Ka1Ka2]/[H⁺]2 (12)

Based on carbonate mass balance and related equilibria, the fractions of different carbonate species are expressed as follows.

f[H2CO3] = [H2CO3]/[CO2(aq)]Total (13)

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From Eq. (3) and Eq. (12),

f[H2CO3] = KHPCO2 /{ KHPCO2{[H⁺]² + Ka1[H⁺] + Ka1Ka2}/[H⁺]²} (14) f[H2CO3] = [H⁺]²/{[H⁺]² + Ka1[H⁺] + Ka1Ka2]} (15) Likewise, it can be shown that

f[HCO3-] = [HCO3-]/[CO2(aq)]Total (16) f[HCO3-] = Ka1[H⁺]/{[H⁺]² + Ka1[H⁺] + Ka1Ka2]} (17) f[CO32-] = [CO32-]/[CO2(aq)]Total = Ka1Ka2/{[H⁺]² + Ka1[H⁺] + Ka1Ka2]} (18)

The Eqs. (16- 18) indicate the concentrations of different carbonate species to be pH dependent. With increase in pH the concentrations of less protonated species will increase.

Of interest is the total carbonate as a function of pH. With increase in pH, the solubility of the atmospheric CO2 in water will increase. Further, Fig. 2 shows the variation in H2CO3, HCO3- and CO32-with pH from which following conclusions are drawn:

(1) Below pH 5, the most of dissolved CO2 exists as H2CO3. (2) Between pH 7 and 10, the dominant species is bicarbonate.

(3) Above pH 11, dissolved CO2 exists largely as CO32-.

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Fig. 2 The fractions of different carbonate species as a function of pH

Source : http://ion.chem.usu.edu/~sbialkow/Classes/3600/Overheads/Carbonate/CO2.html CO2 in Equilibrium with Ocean

Ocean – atmosphere CO2 exchange is one of the most complex natural phenomena. Oceans act as a huge carbon sinks, and store about one- third of CO2 emitted by human activity. The amount of CO2 dissolved in ocean is about fifty times of amount present in the atmosphere. The dissolved carbon dioxide, governed by the reversible equilibria(19-21), is distributed in the form of H2CO3, HCO3-

and CO32-

.. Solar heating divides the ocean surface water from the deep ocean water by a temperature gradient, which is known as thermocline. Thermocline is between 75 to 200 m below the ocean surface. Mixing of the upper warm layer and lower cold layer, which is below the thermocline, is very slow and may take many years. Moreover, the movement of surface water layers to the deep ocean takes centuries. Winds and waves keep the upper layer well mixed.

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Interestingly, as we move away from the Equator towards the poles, the water cools and thermocline disappears near the poles. Cooling has another important consequence that the water near the poles becomes saltier as compared to the water at the Equator. This is due to the exclusion of the salts, when ice forms.

Dissolution of gases in the ocean takes a relatively long time to come to equilibrium.

Equilibration halftime for carbon dioxide air-sea exchange water Eqs.(19-20) is approximately five months. The forward andbackward reactions depend up on pH, temperature and alkalinity.

. CO2(g) + H2O H2CO3(aq) (19)

H2CO3(aq) HCO3-(aq) + H+ (20)

HCO3- CO32-(aq)+ H⁺ (21) The pH of sea water is 8.2, and at this pH the distribution among the three inorganic carbon-containing species, CO2(aq), HCO3–(aq), and CO32–(aq), is about 0.5%, 89%, and 10.5%, respectively. Obviously, the bicarbonate ion, HCO3-, is the predominant inorganic carbon species in the ocean.

The solubility of CO2 in seawater increases with the pressure of the gas(Henry’s law), and decreases as the temperature and salinity of the seawater increase. With increase in temperature, the kinetic energy of gas molecules increases, and, therefore, the tendency of the molecules to go out increases.

Nearly all of the carbon in the Earth system is in its highest oxidation state (+4) in the form of gaseous CO2, carbonate salts, HCO3– and CO32– ions in solution and CaCO3(solid).

Phytoplankton the microscopic organisms reside at the surface of all the oceans. These are important contributors to Earth’s carbon budget. They are photosynthetic organisms.They remove

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CO2 from the ocean-atmosphere system, when they produce the organic molecules for sustaining their life. Many marine organisms use CaCO3 as part of their structure. Corals build some of the most striking of these structures, including coral reefs They use Ca²⁺(aq) and HCO3–(aq) ions present in the surrounding sea to synthesize the intricate CaCO3 scales that adorn their exterior. The chemical reaction involved in the process of making CaCO3, is represented by Eq. (23).

Ca²⁺(aq) + 2HCO3–(aq) CaCO3 + CO2(aq) + H2O. . . (23)

Reaction(23) is also responsible for the formation of sediments. Although CaCO3 is insoluble, the formation of the precipitate at the concentrations and salinity in seawater is slow. Sedimentary rocks, formed from these sediments deep in the Earth constitute the single largest carbon reservoir of about 10 × 10⁸ Giga ton Carbon

Effect of Increase in CO2 on Dissolution in Oceans

As pointed out in a previous section, the amount of CO2 in atmosphere is increasing and so is the partial pressure of CO2. According to Le Chatelier’s principle(see Module for details), increase in atmospheric CO2, which in gas phase, the equilibrium(24) will shift to right and more of CO2 will dissolve in water leading to rise in CO2(aq). Consequently, increase in CO2(aq) will force the equilibrium(25) to move right resulting in the formation of more hydrogen ions and increase in acidity and decrease in pH.

CO2(g) CO2(aq) (24) CO2(aq) + H2O HCO3–

(aq) + H⁺(aq) (25) The increase in CO2(aq)^.would have effect of dissolving CaCO3(Eq. 26). According to Le Chatelier’s principle, an increase in concentration of the product of Eq. (26), i. e., CO2(aq) would force the reaction (26) to move backward resulting in the dissolution of CaCO3.

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Ca²⁺(aq) + 2 HCO3–(aq) CaCO3 + CO2(aq) + H2O (26) The net result of changes in equilibria(24-26) owing to rise in atmospheric CO2 is an increase in overall [H⁺]. It has been experimentally observed that the pH of the top layer of sea water has decreased by 0.1 unit from 8.2 to 8.1 during last hundred years or so. Consequently, the two vital species, CaCO3 and CO32- have become less available due to dissolution of CaCO3 in sea water(Eq.

26). It is estimated that about 100 GIC (equivalent to 370 Gt CO2) has dissolved in the oceans. This portends serious consequences for the survival of shell-forming organisms in an environment that favors dissolution of CaCO3. Since many of these organisms are at the base of the oceanic food chain and are suppliers of significant portion of the human food supply, it is matter of great environmental concern.

At the current rate of rise in CO2, the concentrations of critical species, CaCO3, and CO32-

would continue to decrease and ultimately the effects of all these changes on marine life would become apparent sooner than later. There is a good possibility of shifting of other species, which can better adapt in acidic environment.

Carbonate Solubility and Alkalinity

1. Calcium carbonate solubility in water in a closed system

We consider the solubility of calcite(CaCO3) in water in a closed vessel, not in contact with air.

There shall be two chemical reactions(27-28).

CaCO3(s) Ca²⁺(aq) + CO32-(aq) ;Ksp (27) Carbonate ion, like all weak acid anions, is hydrolyzed.

CO32-(aq)+ H2O HCO3- (aq)+ OH^- ; Kh or Kb ( 28)

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Addition of Eqs. (27-28), we obtain:

CaCO3(s) + H2O Ca²⁺(aq) + HCO3-(aq) + OH^- ; K (29) Equilibrium constants for Eqs.(27 – 29) are defined by Eqs.(30=32). Ksp is the solubility product of CaCO3, Kh is hydrolysis constant(also known as base dissociation constant, Kb) of carbonate ion and K represents the equilibrium constant of the reaction(29).

Ksp = [Ca²⁺][CO32-] (30)

Kh = {HCO3-(aq)][OH^-]/[CO32-(aq)] (31) K = [Ca²⁺(aq)][HCO3-(aq)][OH^-] (32) On multiplying Eqs.(30-31), we get:

KspKh = [Ca²⁺][HCO3-(aq)][OH^-] (33)

On comparing Eqs.(32) and (33)we find:

K = KspKh = [Ca²⁺][HCO3-(aq)][OH^-] (34)

Using the literature values of Ksp = 10^-84 and Kh = 10^-37 at 25^oC, from Eq. (34), we get, K = 10^-

12.01.

Remembering that in CaCO3, the molar concentrations, [Ca²⁺] = [carbonate ion]. Since HCO3- is formed from CO32-, by mass balance:

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[Ca²⁺] = [CO32-] + [HCO3-] (35) From Eq.(29)

[HCO3-] = [OH^-] (36)

On substituting the values of HCO3- from Eq. (36) in Eq. (34):

K = [Ca²⁺][HCO3-(aq)][ HCO3-(aq)] = = [Ca²⁺][HCO3-(aq)]² (37) In terms of calcium ion concentration,

[HCO3-(aq)] = { K/ [Ca²⁺]}^1/2 (38) From Eq. (30):

[CO32-] = Ksp/[Ca²⁺] (39) On substituting the values of [HCO3-(aq)] and [CO32-] from Eqs.(38) and (39), respectively in equality(35), we obtain:

[Ca²⁺]² = [CO32-] + { K/ [Ca²⁺]}^1/2 (40) An algebraic exact solution of the Eq. (40) is difficult, but by successive approximations it can be solved for [Ca²⁺] and subsequently for [OH^-]. The latter allows the calculation of pOH and then pH. Finally, this procedure yields a pH value of 9.94 for a saturated solution of CaCO3

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2. Carbonate Solubility in a System Open to Air

It must be remembered that CO2 when it dissolves in water, it reacts with CO32- ; there is no reaction between CaCO3(solid) and gaseous CO2. In an open system in equilibrium with atmosphere, the major reactions are Eqs. (41) and (42).

CaCO3(s) Ca²⁺(aq) + CO32-(aq) (41) CO32- + CO2(aq){ or H2CO3 } 2HCO3- (42)

Equation(42) is the sum of the reactions(28) and (43). The OH^-formed in Eq. (28) reacts with CO2(Aq)( or H2CO3) as in Eq. (43)

CO32-(aq)+ H2O HCO3- (aq)+ OH- (28)

CO2(Aq)( or H2CO3) + OH- HCO3- + H2O (43)

On combining Eqs. (41 and (42), we get Eq. (44).

CaCO3(s) + CO2(Aq)( or H2CO3) Ca²⁺(aq) + 2HCO32-(aq) (44) The equilibrium constant for reaction(44), Kopen , is defined by Eq. (45).

Kopen = [Ca²⁺] [HCO3-(aq)]² / [H2CO3] (45).

Kopen = KspKa1/ Ka2

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Using Ksp = 10^-8.34; Ka1 = 10^-6.4; Ka2 = 10^-10.33, we obtain, Kopen = 10^-4.41.

Stoichiometric Eq. (44), shows that the molar concentration of HCO3- formed is twice the molar concentration of Ca^2+,, it is obvious therefore that

[HCO3-] = 2 [Ca²⁺] (46) On substituting the value of [HCO3-] from Eq. (46) into Eq. (45), we obtain:

Kopen = [Ca²⁺] ×4[Ca²⁺]² / [H2CO3] = 4[Ca²⁺]³ / [H2CO3] (47)

[Ca²⁺] = { Kopen[H2CO3]/4}^1/3 (48)

According to Henry’s law, at fixed atmospheric PCO2 , the value of dissolved CO2 in the form of [H2CO3] shall be fixed(Eq. 19).

At 370 ppm atmospheric CO2, PCO2 is 370 × 10^-6 atm. Using KH = 10^-1.5 mol L^-1 atm, we calculate;

[H2CO3] = KHPCO2 = 370 × 10^-6 (atm)× 10^-1.5( mol L^-1 atm) = 10^-4.9mol L^-1.

On substituting this value of [H2CO3] and Kopen value in Eq. 47, we obtain: [Ca²⁺] = 10^-3.3. From the relation(46):

[HCO3-] = 2 [Ca²⁺] = 2×10^-3..3 = 10^-3.0 mol L^-1. (49) To calculate pH, we use the equations( 50) and (51).

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Ka1

H2CO3- H⁺+ HCO3- (50) Ka1 = [H⁺] [HCO3-]/[H2CO3] (51) Eq. (51) on rearrangement becomes(Eq. 52).

[H⁺] = Ka1[H2CO3]/[HCO3-] (52) The logarithmic form of Eq. (52) is:

log[H⁺] = log Ka1 + log[H2CO3]/[HCO3-] (53)

Multiplying both sides of Eq. (53) by -1, we obtain:

-log[H⁺] = -log Ka1- log[H2CO3]/[HCO3-] (54) Replacing -log[H⁺] by pH and -log Ka1 by pKa1 (see Module 5) we get:

pH = pKa1 - log[H2CO3]/[HCO3-] (55) As shown earlier, pKa1 = 6.4, [H2CO3] = 10^-4.9mol L^-1 and [HCO3-] = 2[Ca²⁺] = 10^-3 mol L^-1, Eq. (56) yields:

pH = 6.4 – log{10^-4.9/10^-3} (56) pH = 8.3

So the pH of an open aqueous system in equilibrium with atmospheric CO2 is lower by about 1.64 unit. It can be shown that the solubility of CaCO3 is about four times more in an open system than in a closed system.

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Alkalinity

The capacity of a water sample to accept hydrogen ions( H⁺) to neutralize acids is called alkalinity. Total alkalinity is the amount of acid needed to bring the sample pH to 4.5. At this pH, all the alkaline compounds in the sample are used up.’ A knowledge of the alkalinity of waters of different types is important from the view point of understanding the chemistry and biology of waters.

For use in industries, in food processing, for in water for drinking, etc., it is necessary to know the level of alkalinity. This knowledge helps calculate the amount of alkali neutralizing chemicals to be added to neutralize the alkaline materials present in water to be used. If the wastewater released by any unit in a stream is acidic, it will immediately lead to acidification of river if the river does not have sufficient acid neutralizing capacity. Thus, it is always necessary to know the alkalinity of the stream before releasing effluent into it.

In natural waters, in general the species responsible for alkalinity are bicarbonate, carbonate and hydroxide ions. In some cases, other basic substances such as ammonia, phosphate, borate, silicate, acetate ions, etc may also be present. The alkaline species neutralize the acid as in Eqs. (57- 59).

HCO3- + H⁺ CO2 + H2O (57) CO32- + H⁺ HCO3- (58)

OH^- + H⁺ H2O (59)

Reactions(57 – 59) go almost to completion and hence are not written as equilibria using equilibrium sign. The total alkalinity of a water sample is expressed by Eq. (60).

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[Total Alkalinity] = [HCO3-] + 2[CO32-] + [OH^-] – [H⁺] (60)

Since carbonate for complete neutralization needs two hydrogen ions, [CO32-] has been multiplied by two. At pH more than 7, the last term on right hand side of the Eq. (60) becomes negligible and then Eq. (60) modifies to Eq. (61)

[Total Alkalinity] = [HCO3-] + 2[CO32-] + [OH^-] (61) Alkalinity has two scales.

1. Phenolphthalein Alkalinity: This value is corresponding to titration with an acid(H2SO4) to the phenolphthalein indicator end point pH 8.3, where predominant carbonate species is HCO3- ..

2. Total Alkalinity: This value corresponds to the titration with an acid(H2SO4) to the methyl orange end point pH 4.5.

Experimental details of Alkalinity determination are in web sites cited.

Suggested Reading

1. T. G. Spiro and William M. Stigliani(2003), Chemistry of the Environment, Prentice –Hall of India, New Delhi.

2. P. V. Hobbs(2000), Basic Physical Chemistry for Environmental Sciences, Cambridge, UK 3. American Chemical Society (2015), Ocean Chemistry; ACS Climate Science Toolkit,

Washington DC

4. Chen-Tung Arthur Chen, Carbonate Chemistry of Oceans, Oceanography, Vol. 1, Encyclopedia of Life Support Systems, seen on web.

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5. http://nitttrc.ac.in/four%20quadrant/eel/quadrant%20-%201/exp7_pdf.pdf, Seen on web on August, 20, 2015.

6. P.V. Hobbs(2000), Introduction to Atmospheric Chemistry, Cambridge University Press, Cambridge

7. P. Brimblecombe(1996) Air Composition and Chemistry, Cambridge University Press, Cambridge.

8. Colin Baird (1998), Environmental Chemistry, W.H. Freeman, New York.

9. R. F. Phelan and R. N. Phelan(2013), Introduction to Air Pollution Chemistry- A Public Health Perspective, Jones and Bartlett, Burlington, MA.

10. J. Gerard(2012), Introduction to Environmental Chemistry, Jones and Bartlett, New Delhi 11. file:///C:/Users/PROF.K.S.GUPTA/Desktop/carbonate%20system/Carbon%20Dioxide%20-

%20Carbonic%20Acid%20Equilibrium.html

12. file:///C:/Users/PROF.K.S.GUPTA/Desktop/carbonate%20system/Carbon%20dioxide%20-

%20Wikipedia,%20the%20free%20encyclopedia.html

13. file:///C:/Users/PROF.K.S.GUPTA/Desktop/carbonate%20system/Carbon%20Dioxide%20and

%20Carbonic%20Acid.html