#### Chemical Parameters in Water Chemistry and Analysis, Chemistry tutorial

Introduction

Like physical parameters, several chemical parameters are paramount to crucial decision-making concerning the quality of public water supplies. Among these are pH, acidity, alkalinity and hardness.

Some Relevant Chemical Parameters of Water

pH

pH is a term utilized to state the intensity of the acid or alkaline situation of a solution. It is a way of expressing the hydrogen-ion concentration or more precisely, the hydrogen-ion activity. It is significant in approximately every phase of environmental science and engineering. In the field of water supplies, it is a factor that must be considered in chemical coagulation, disinfection, water softening and correction control. In wastewater treatment utilizing biological procedures, pH must be controlled inside a range favorable to the particular organisms engaged. Chemical processes utilized to coagulate wastewaters, dewater sludges or oxidise indeed substance, these as cyanide ion, need that the pH be controlled inside rather narrow limits.

pH and the Ionic Product of Water (KW)

The pH of a solution is described as the negative logarithm to base ten of the hydrogen ion concentration of that solution. That is,

pH = - log10 [H+] or pH = log10 1/ [H+]

The pH scale ranges from O to 14, through pH 7 at 250C representing absolute neutrality. Acid range Neutrality Alkaline range

At 2500C, pure water dissociates as follows:

H2   H+   + OH-

The [H+], using the hydrogen electrode, of pure water at 250C has been established to be 10-7? This implies that the [OH-] will also be 10-7. The equilibrium equation provides: [H+][OH]

K = [H2 O]

Since the ionisation degree is so small and the concentration is so huge, it is considered that the concentration of water is steady. Therefore,

Kw = [H+] [OH-]

For pure water at 250C, Kw = 10-7 x 10-7 = 10-14

By implication, pH = 7 and pOH =7 So, pH + pOH = 14

Acidity

Acidity is calculated of the ability of a specified water example to neutralize strong bases to an indicator end point. Most natural waters, domestic wastewaters and many industrial wastes are buffered principally via a carbon dioxide-bicarbonate system.

Sources and Nature of Acidity

Carbonic and species are shaped when CO2 enters surface waters. This occurs when the concentration of CO2 in water is less than that in equilibrium through CO2 in the environment. Through biological oxidation of organic matter chiefly in polluted water, CO2 created can too be absorbed via water.

Ground waters and waters from the hypolimnion of stratified lakes and reservoirs frequently include considerable amounts of CO2 resulting from bacterial oxidation of organic matter with which the water has been in contact. Under these conditions, the CO2 isn't free to escape to the atmosphere. Carbon dioxide is an end product of both aerobic and anaerobic bacteria oxidation; therefore, its concentration isn't bounded via the amount of melted oxygen originally present. It isn't uncommon to encounter ground waters through 30 to 50 mg/L of CO2.

Mineral acidity is present in many industrial wastes chiefly those of the metallurgical industry and several from the production of synthetic organic materials. The drainage from abandoned mines and iron ore dumps will enclose significant amounts of sulphuric acid or salts of sulphuric acid if sulphure, sulphide or iron pyrites are present.

Conversion of such substances to sulphuric acid and sulphate is brought about via sulphure-oxidizing bacteria under aerobic conditions.

2S + 3O2 + 2H2O bact.  4H+ + 2SO42-

FeS2 + 3½ O2 + H2O bact Fe2+ + 2H++2SO42-

Salts of trivalent heavy metals particularly for example Fe3+ and Al3+ hydrolyse in water to liberate mineral acidity.

FeCl3+ 3H2O      Fe (OH)3 + 3H+ + 3Cl-

Many industrial wastes enclose organic acids. Combination of fossil fuels in power plants and automobiles directs to the arrangement of oxides of nitrogen and sulphure, which when mixed with rain, hydrolyse to form sulphuric and nitric acids.

Significance of Acidity

Acid waters are of concern because of their corrosive characteristics and the expenses involved in removing or controlling the corrosion- producing substances. The corrosive factor in most waters is CO2, but in many industrial wastes, it is mineral acidity. Acid rain can lower the pH in poorly buffered lakes thereby adversely affecting aquatic life, and can increase the amount of chemical, such as aluminum, leached from soil into surface.

Methods of Acidity Measurement

Both CO2 and mineral acidity can be measured by means of standard solutions of alkaline reagents. Mineral acids are measured by titration to a pH of about 3.7, the methyl orange end point. For this reason, mineral acidity is also called methyl orange acidity. Titration of a example to the phenolphthalein end point of pH 8.3 measures both mineral acidity plus acidity due to weak acids. This total acidity is also termed phenolphthalein acidity.

Determination of Acidity in Natural and Waste Waters

(a) Methyl Orange Acidity: While methyl orange was formerly utilized for this purpose, bromophenol blue is now recommended as it has a sharper colour change at pH 3.7. The titration is carried out using 0.02M NaOH. Consequences are reported in terms of methyl orange acidity expressed as CaCO3. That is,

Acidity (as mg/L CaCO3) = where V = mL sodium hydroxide titrant.

M = molarity of sodium hydroxide VxMx100000 mLsample

The molecular weight of CaCO3 = 100g (=100,000mg)

(b) Phenolphthalein acidity:  This measures the total acidity resulting from both mineral acids and weak acids in the sample. Either phenolphthalein or metacresol purple sign can be utilized for this titration. When heavy-metal salts are present, it is usually desirable to heat the sample to boiling and then carry out the titration. The heat speeds the hydrolysis of the metal salts, permitting the titration to be completed more readily. Again, 0.02M NaOH is utilized as the titrating agent. Consequence are reported in terms of phenolphthalein acidity stated as CaCO3 as before.

Determination of Total Acidity via Mixed Indicators technique the mixed indicator is prepared via mixing 10mL of 0.1percent thymol blue (in 50percent ethanol) through 30mL of 0.1percent phenolphthalein (in 50percent ethanol). Compute precisely 50 or 100mL of water example into a titration flask. Add one drop of the mixed indicator. Titrate through 0.0125M barium hydroxide solution to the end point yellow (acid) to violet noticed by a pH meter rather than via gauge.

Acidity (as mg/L CaCO2) = VxMx100000 mLsample

where V = Ml barium hydroxide

M = molarity of barium hydroxide

100000 = molar weight of CaCO3 (for example 100g).

Alkalinity

The alkalinity of a water sample is a measure of its capacity, to neutralize acids. Constituents of alkalinity in natural water systems comprise CO32-, HCO3-, OH-, HSiO3-, H2BO3-, HPO4-, HS-   and NH30. Such compounds consequence from the dissolution of mineral materials in the soil and atmosphere. Phosphates may also originate from detergents in waste water discharges and from fertilizers and insecticides from agricultural land. Hydrogen sulphide and ammonia may be products of microbial decomposition of organic material.

By far, the most common constituents of alkalinity are bicarbonate (HCO3-), carbonate (CO32-) and hydroxide (OH-). In addition to their mineral origin, these substances can originate from CO2, a constituent of the atmosphere and a product of microbial decomposition of organic material, according to the following reactions.

CO2 + H2  H2CO3 (dissolved CO2 and carbonic acid)

H2CO3H+ + HCO3- (bicarbonate)

HCO-3     H+ + HCO32- (carbonate)

CO2- H2O   HCO3-   + OH- (hydroxide).

The last reaction is a weak reaction chemically. Though, utilisation of the bicarbonate ion as a carbon source via algae can drive the reaction to the right and consequence in substantial accumulation of OH-. Water through heavy algal growths frequently has pH values as high as 9 to 10.

Significance and Application of Alkalinity Data

The principal objection to alkaline water is the reactions that can take place between alkalinity and indeed actions in the water. The consequential precipitate can foul pipes and other water-system appurtenances. Alkalinity is a significant consideration in calculating the lime and soda-ash requirements in softening of water via precipitation process. It is as well a means of evaluating the buffering capacity of waste waters and sludges.

Determination of Alkalinity in Natural and Treated Waters (Titration Method)

In the natural and treated waters alkalinity determination, four quantities are usually reported.  Such are phenolphthalein alkalinity, total alkalinity, carbonate alkalinity and total carbon dioxide.

Determination of Phenolphthalein Alkilinity (PA)

Put 50 mL or 100 mL of water sample into a clean conical flask. Add one drop of 0.05M sodium throsulphate solution to eliminate free residual chlorine if present. Add 2 drops of phenolphthalein indicator. If the solution continues coloured, then PA = 0. If the solution turns red, PA is present.  Then  titrate  the  solution  with  0.02M  HCl  until  the  colour disappears.

The conversion attained in this titration corresponds to:

OH- + CO2-3+ 2H+     H2O + HCO - 3

Now,

PA (as mg/L CaCO3) =   VP xMx100000 /mLsample

where Vp = volume (mL) of the acid used

M = molarity of the acid

Determination of Total Alkalinity (TA)

Add 2 drops of blended indicator (bromocresol green + methyl red solution) or of methyl orange indicator into 50 or 100 mL of water example in a clean conical flask. Shake and titrate through 0.02M HCl acid until, at pH 4.6, the colour transforms to pink (for mixed indicator) or from yellow to orange (for methyl orange indicator). The conversion for this titration corresponds to

OH- + CO2-3 + HCO -3 + 4H+     3H2O + 2CO2

Now,

TA (as mg/L CaCO3) = VT xMx100000 /mLsample

Where VT = volume (mL) of acid utilized

M = molarity of acid utilized.

Determination of Hydroxide, Hydrogen Carbonate and Carbonate to 50 or 100 mL of the water sample in a clean volumetric flask, add a slight excess of BaCl2 solution to precipitate the carbonate. The HCO3- and OH- aren't affected; the HCO- is not also affected by phenolphthalein indicator. Add two drops of phenolphthalein indicator and titrate the OH- in the water sample against 0.02M HCl until the solution is colourless. Let the volume of the acid used be VH mL. Add two drops of the mixed indicator or methyl orange to the solution of the hydroxide/ acid titration. Shake and titrate the HCO3- to the end point with 0.02M HCl. Let the volume of acid used be VHC mL.

To a fresh 50 or 100 mL water sample, add 2 drops of mixed indicator or methyl orange and shake. Titrate to the end point through 0.02M HCl. The volume of the acid used, VT, is for the three species OH-, HCO3- and CO32-. Hence, the volume of acid utilized for CO32- only, VC = [VT (VH + VHC)] mL. The alkalinities can now be computed as usual.

Determination of Total CO2

Collect the water sample into a 500 mL flask leaving no air space. Take to the laboratory as soon as possible and siphon into a 100 mL graduated cylinder allowing overflow to take place. Add five to ten drops of phenolphthalein indicator. If the water example turns red, the free-CO2 is absent, but if the water continues colourless, titrates speedily through a standard Na2CO3 or NaOH solution until pink colour persists for about 30 seconds.

Alkalinity as mg/L CO2 = VxMx44000/ mLwatersample

where

V= volume (mL) of Na2CO3 OR NaOH utilized

M = molarity of the alkali (Na2CO3 or NaOH)

44000 = molar weight of CO2 in mg.

Hardness

Hardness is described as the concentration of multivalent metallic cations in water which find out the capacity of the water to precipitate soap. Depending on the anion with that it associates, hardness is classified as calcium and magnesium hardness, carbonate hardness and noncarbonate hardness and pseudo-hardness. Carbonate hardness is sensitive to heat and precipitates readily at high temperatures, for example

Ca (HCO3)2      CaCO3 + CO2 + H2O

Water hardness is due mainly to the presence of Ca2+ and Mg2+ in water. Other ions that may cause hardness include Fe2+, Mn2+, Sr2+ and Al3+. The latter are originate in much smaller quantities than Ca2+ and Mg2+, and for all practical purposes, hardness might be symbolized via the sum of the Ca2+ and Mg2+ ions in a specified water example.

Impacts of Water Hardness

Soap consumption via hard waters symbolizes an economic loss to the water user. Sodium soaps react through multivalent metallic cations to form a precipitate, thereby losing their surfactant properties.

2NaCO2 C17H35 + Ca2+   Ca2+(CO2  C17H35)2  + 2Na+

Soap precipitate

Lathering doesn't take place until all of the hardness ions are precipitated. The precipitate shaped adheres to surfaces of tubs, sinks, dishwashers and might stain clothing, dishes and other items.  Residues of the precipitate might continue in the pores so that skin might feel rough and uncomfortable. Boiler scale, consequential from carbonate hardness, may reason considerable economic loss through fouling of water heaters and hot-water pipes. Changes in pH of the water distribution systems may as well consequence in deposits of precipitates.  Bicarbonates start to convert to the less soluble carbonates at pH values above 9.0.

Magnesium hardness, particularly connected through the sulphate ion, has a laxative result on persons unaccustomed to it. Magnesium concentrations of less than 50 mg/L are desirable in potable waters even though many public water supplies exceed this amount.

Determination of Hardness

Hardness can be computed via using spectrophotometric techniques or chemical titration to determine the quantity of calcium and magnesium ions in a specified water sample.  Hardness can be computed directly through titration by ethylenediamine tetraacetic acid (EDTA) using Eriochrome Black T (EBT) as an indicator. The EBT reacts by the divalent metallic cations, shaping a compound, which is red in colour. The EDTA swaps the EBT in the compound, and when the replacement is complete, the solution transforms from red to blue.

Mn+ + EBT M - EBT

M - EBT + EDTA M - EDTA + EBT

Red                                blue

If 0.01M EDTA is utilized, 1.0 mL of the titrant measures 1.0 mg of hardness as CaCO3.

Application of Hardness Data

Hardness of water is a significant consideration in determining the suitability of a specified water source for domestic and industrial utilizes. The environmental engineer utilizes it as a basis for recommending the need for softening procedures and design kinds. Hardness might range from virtually zero to numerous hundred or thousand parts per million. Even though acceptability levels fluctuate according to a consumer's acclimation to hardness, a usually allowed classification is as follows:

Soft water < 50 mg/L as CaCO3

Moderately hard water 50 - 150 mg/L as CaCO3

Hard water 150 - 300 mg/L as CaCO3

Very hard water > 300 mg/L as CaCO3

The Public Health Service Standards recommend a maximum of 500 mg/L of hardness in drinking water.

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