Mineral Water Quality Inspection

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SAMPLE QUALITY

             We analyses the samples of drinking water and perform tests which conclude that all samples are differs in their characteristic properties but finally qualifies as drinking water. These Properties create difference in quality of water. A brief description on each prperty is given below:

pH

S.No	Water Sample	Observed pH	Actual pH
1.	Nestle	8.41	6.5 – 8.5
2.	Spring	6.42	7 – 3
3.	Aquafina	8.28	6.5 – 8.5

pH value determines whether water is hard or soft. The pH of pure water is 7.The normal range for pH in surface water systems is 6.5 to 8.5 and for groundwater systems 6 to 8.5.water with a low pH (< 6.5) could be acidic, soft, and corrosive. Therefore, the water could contain metal ions such as iron, manganese, copper, lead, and zinc...or, on other words, elevated levels of toxic metals. This can cause premature damage to metal piping, and have associated aesthetic problems such as a metallic or sour taste, staining of laundry, and the characteristic "blue-green" staining of sinks and drains. More importantly, there are health risks associated with these toxins. Water with a pH > 8.5 could indicate that the water is hard. Hard water does not pose a health risk, but can cause aesthetic problems. These problems include an alkali taste to the water (making that morning coffee taste bitter!), formation of a deposit on dishes, utensils, and laundry basins, difficulty in getting soaps and detergents to lather, and formation of insoluble precipitates on clothing.

Result
The ph range of nestle and aquafina is in the required limit while the ph of spring is not in the required limit.

Chloride Content

S.No	Water Sample	Observed Chloride content (mg Cl / L)	Actual Chloride content (mg Cl / L)
1.	Nestle	54.98	77 - 150
2.	Spring	29.9	38
3.	Aquafina	4.99	20

chloride presents a pretty good problem in large concentrations, giving water that briny, brackish taste that would make anybody's tongue cringe in distaste. Usually chloride concentrations found in water supplies are low, which is a good thing since low to moderate chloride concentrations can add a pleasant taste to your drinking water, making chloride desirable for this particular reason, while high concentrations of chloride ions can add to the electrical conductivity of the water EPA Secondary Drinking Water Regulations recommended the maximum concentration of 250mg/1 for chloride ions.we don’t need as much chloride because of the regular intake of Nacl so chloride content should be less in the water.

Result

On the basis of above theory all the chloride value of the sample is in the acceptable range. But it is good that chloride content should be in moderate range to have a good taste so spring is good as compare to others.

Hardness

Serial number	Sample	Observed Calcium hardness (mg CaCO3 / L)	ActualCalcium hardness (mg CaCO3 / L)
1.	Nestle	120	40 - 70
2.	Spring	50	Not mention
3.	Aquafina	120	54

              Hard water is not a health hazard. In fact, the National Research Council (National Academy of Sciences) states that hard drinking water generally contributes a small amount toward total calcium and magnesium human dietary needs. They further state that in some instances, where dissolved calcium and magnesium are very high, water could be a major contributor of calcium and magnesium to the diet.

Result

On the basis of above we can conclude that nestle is good because of high range of hardness.

                                   Total Dissolved Solid

S.No	Samples	Observed T.D.S. mg/L	Actual T.D.S. mg/L
1.	Nestle	110	160 - 350
2.	Spring	120	259.7
3.	Aquafina	120	105

The electrically charged dissolved particles make ordinary natural water a good conductor of electricity. Coversely, pure water has a high electrical resistance, and resistance is frequently used as a measure of its purity.

Since only a few of these most common ionic water contaminants are health related, most natural water supplies are safe to drink from the standpoint of dissolved inorganic chemical contaminants. However, even though found more rarely -- and in much smaller quantities -- certain inorganic ions can be toxic.

Result

So on the basis of above description it can be concluded that we can be concluded which one is better in them after knowing the complete information of these electric charge dissolved in it. But to have a low limit of tds is so aquafina is good because of having low tds content.

Sulphate Content

S.No	Sample	Observed Sulphate Content mg SO4/L	Actual Sulphate Content mg SO4/L
1.	Nestle	11.6	7-30
2.	Spring	11.5	59
3.	Aquafina	12.3	52

Ø  High amounts of various sulphate salts may give drinking water an offensive taste. Depending upon the type of sulphate salt present in the water, most people begin to notice an offensive taste at concentrations ranging from 250 to 1,000 mg/L.

Ø  High concentrations of sulphate may interfere in the efficiency of chlorination in some water supplies.

Ø  Also, sulphate salts may increase the corrosive properties of water.

Ø  Air may contain sulfates in areas of heavy industry. Many sulfate salts can react in air to form dilute acid, which can irritate intestinal pain

Ø  dehydration as a result of diarrhea

Ø slight decrease in normal stomach acidic nestle is best because of low content of sulfate

Result

Nestle is best because of low content of sulfate

Alkalinity

S.No	Sample	Observed Total Alkalinity (meq/Lit)
1.	Nestle	60
2.	Spring	150
3.	Aquafina	60

Alkalinity of water may be due to the presence of one or more of a number of ions. These include hydroxides, carbonates and bicarbonates. Small amounts of carbonates are found in natural water supplies in certain sections of the country, rarely exceeding 3 or 4 gpg. Bicarbonates are the most common sources of alkalinity. Almost all natural supplies have a measurable amount of this ion, ranging from 0 to about 50 gpg. Phosphates and silicates are rarely found in natural supplies in concentrations significant in the home. Moderate concentrations of alkalinity are desirable in most water supplies to balance the corrosive effects of acidity. However, excesive quantities cause a number of problems. These ions are, of course, free in the water, but have their counterpart in cations such as calcium, magnesium and sodium or potassium.You probably will not notice an alkaline condition due to bicarbonate ions except when present in large amounts. In contrast, you should readily detect alkalinity due even to fairly small amounts of carbonate and hydroxide ions.

Strongly alkaline Waters have an objectionable "soda" taste. The EPA Secondary Drinking Water Regulations limit alkalinity only in terms of total dissolved solids (500 ppm) and to some extent by the limitation on pH.

Highly mineralized alkaline waters also cause excessive drying of the skin due to the fact that they tend to remove normal skin oils.

 RESULT

On the basis of description we can conclude that the spring and aquafina are comparativele good compare to nestle because nestle contain high alkalinity.

The Overall Results Of Water Analysis:

water treatment	ph	Chlorides (mg/L)	Hardness (mg/L)	TDS (mg/L)	Alkanity (mg/L)	Sulphate (mg/L)
Nestle	8.41	54.98	120	110	60	11.6
Spring	6.42	29.9	50	120	150	11.5
Aquafina	8.28	4.99	120	120	60	12.3

Sources Of Error

There are several sources of error due to which our results are deviated from the standards result

The apparatus are calliberated at 25 centigrade

For washing we use distilled water which already contain many impurities

Apparatus is not properly washed

Any chromic acid content remain after washing causing the error in result

The concentration of reagent may be changed due to prolonged storage

Some time colour changed was not sharped they can also effect on the readind

Human error

DISCUSSION

pH:

pH is a measure of the acidity or basicity of a solution. It is defined as the cologarithm of the activity of dissolved hydrogen ions (H⁺). Hydrogen ion activity coefficients cannot be measured experimentally, so they are based on theoretical calculations. The pH scale is not an absolute scale; it is relative to a set of standard solutions whose pH is established by international agreement.
The concept of pH was first introduced by Danish chemist Søren Peder Lauritz Sørensen at the Carlsberg Laboratory in 1909. It is unknown what the exact definition of p is. Some references suggest the p stands for “Power”, others refer to the German word “Potenz” (meaning power in German), still others refer to “potential”. Jens Norby published a paper in 2000 arguing that p is a constant and stands for “negative logarithm”; which has also been used in other works. H stands for Hydrogen. Sørensen suggested the notation "PH" for convenience, standing for "power of hydrogen", using the cologarithm of the concentration of hydrogen ions in solution, p[H] Although this definition has been superseded p[H] can be measured if an electrode is calibrated with solution of known hydrogen ion concentration.
Pure water is said to be neutral. The pH for pure water at 25 °C (77 °F) is close to 7.0. Solutions with a pH less than 7 are said to be acidic and solutions with a pH greater than 7 are said to be basic or alkaline. pH measurements are important in medicine, biology, chemistry, food science, environmental science, oceanography and many other applications.
pH is defined as minus the decimal logarithm of the hydrogen ion activity in an aqueous solution. By virtue of its logarithmic nature, pH is a dimensionless quantity.

pOH

pOH is sometimes used as a measure of the concentration of hydroxide ions, OH⁻, or alkalinity. pOH is not measured independently, but is derived from pH. The concentration of hydroxide ions in water is related to the concentration of hydrogen ions by                                       [OH⁻] = K_W /[H⁺]
where K_W is the self-ionisation constant of water. Taking cologarithms

pOH = pK_W − pH.

So, at room temperature pOH ≈ 14 − pH. However this relationship is not strictly valid in other circumstances, such as in measurements of soil alkalinity.

Total dissolved solids:

Total Dissolved Solids (often abbreviated TDS) is an expression for the combined content of all inorganic and organic substances contained in a liquid which are present in a molecular, ionized or micro-granular (colloidal sol) suspended form. Generally the operational definition is that the solids must be small enough to survive filtration through a sieve size of two micrometres. Total dissolved solids are normally only discussed for freshwater systems, since salinity comprises some of the ions constituting the definition of TDS. The principal application of TDS is in the study of water quality for streams, rivers and lakes, although TDS is generally considered not as a primary pollutant (e.g. it is not deemed to be associated with health effects), but it is rather used as an indication of aesthetic characteristics of drinking water and as an aggregate indicator of presence of a broad array of chemical contaminants.
Primary sources for TDS in receiving waters are agricultural runoff, leaching of soil contamination and point source water pollution discharge from industrial or sewage treatment plants. The most common chemical constituents are calcium, phosphates, nitrates, sodium, potassium and chloride, which are found in nutrient runoff, general stormwater runoff and runoff from snowy climates where road de-icing salts are applied. The chemicals may be cations, anions, molecules or agglomerations on the order of 1000 or fewer molecules, so long as a soluble micro-granule is formed. More exotic and harmful elements of TDS are pesticides arising from surface runoff. Certain naturally occurring total dissolved solids arise from the weathering and dissolution of rocks and soils. The United States has established a secondary water quality standard of 500 mg/l to provide for palatability of drinking water.
Total dissolved solids are differentiated from total suspended solids (TSS), in that the latter cannot pass through a sieve of two micrometres and yet are indefinitely suspended in solution. The term "settleable solids" refers to material of any size that will not remain suspended or dissolved in a holding tank not subject to motion, and exclude both TDS and TSS. Settleable solids may include larger particulate matter or insoluble molecules.

Measurement of TDS:

The two principal methods of measuring total dissolved solids are gravimetry and electrical conductivity. Gravimetric methods are the most accurate and involve evaporating the liquid solvent to leave a residue which can subsequently be weighed with a precision analytical balance (normally capable of .0001 gram accuracy). This method is generally the best, although it is time consuming and leads to inaccuracies if a high proportion of the TDS consists of low boiling point organic chemicals, which will evaporate along with the water. In the most common circumstances inorganic salts comprise the great majority of TDS, and gravimetric methods are appropriate.


Electrical conductivity of water is directly related to the concentration of dissolved ionized solids in the water. Ions from the dissolved solids in water create the ability for that water to conduct an electrical current, which can be measured using a conventional conductivity meter or TDS meter. When correlated with laboratory TDS measurements, electrical conductivity provides an approximate value for the TDS concentration, usually to within ten percent accuracy.

Hard water:

Hard water is water that has high mineral content (mainly calcium and magnesium ions) (in contrast with soft water). Hard water minerals primarily consist of calcium (Ca²⁺), and magnesium (Mg²⁺) metal cations, and sometimes other dissolved compounds such as bicarbonates and sulfates. Calcium usually enters the water as either calcium carbonate (CaCO₃), in the form of limestone and chalk, or calcium sulfate (CaSO₄), in the form of other mineral deposits. The predominant source of magnesium is dolomite (CaMg(CO₃)₂). Hard water is generally not harmful to one's health.
The simplest way to determine the hardness of water is the lather/froth test: soap or toothpaste, when agitated, lathers easily in soft water but not in hard water. More exact measurements of hardness can be obtained through a wet titration. The total water 'hardness' (including both Ca²⁺ and Mg²⁺ ions) is read as parts per million (ppm) or weight/volume (mg/L) of calcium carbonate (CaCO₃) in the water. Although water hardness usually only measures the total concentrations of calcium and magnesium (the two most prevalent, divalent metal ions), iron, aluminium, and manganese may also be present at elevated levels in some geographical locations.

Types of hard water

In the 1960s, scientist Chris Gilby Jnr discovered that hard water can be categorized by the ions found in the water.^{[citation needed]} A distinction is also made between 'temporary' and 'permanent' hard water.

Temporary hardness

Temporary hardness is caused by a combination of calcium ions and bicarbonate ions in the water. It can be removed by boiling the water or by the addition of lime (calcium hydroxide). Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.
The following is the equilibrium reaction when calcium carbonate (CaCO₃) is dissolved in water:
CaCO₃(s) + H₂CO₃(aq) ⇋ Ca²⁺(aq) + 2HCO₃^-(aq)
Upon heating, less CO₂ is able to dissolve into the water (see Solubility). Since there is not enough CO₂ around, the reaction cannot proceed from left to right, and therefore the CaCO₃ will not dissolve as rapidly. Instead, the reaction is forced to the left (i.e. products to reactants) to re-establish equilibrium, and solid CaCO₃ is formed. Boiling the water will remove hardness as long as the solid CaCO₃ that precipitates out is removed. After cooling, if enough time passes the water will pick up CO₂ from the air and the reaction will again proceed from left to right, allowing the CaCO₃ to "re-dissolve" into the water.
For more information on the solubility of calcium carbonate in water and how it is affected by atmospheric carbon dioxide, see calcium carbonate.

 Permanent hardness

Permanent hardness is hardness (mineral content) that cannot be removed by boiling. It is usually caused by the presence of calcium and magnesium sulfates and/or chlorides in the water, which become more soluble as the temperature rises. Despite the name, permanent hardness can be removed using a water softener or ion exchange column, where the calcium and magnesium ions are exchanged with the sodium ions in the column.
Hard water causes scaling, which is the left over mineral deposits that are formed after the hard water had evaporated. This is also known as limescale. The scale can clog pipes, ruin water heaters, coat the insides of tea and coffee pots, and decrease the life of toilet flushing units.
Similarly, insoluble salt residues that remain in hair after shampooing with hard water tend to leave hair rougher and harder to untangle.
In industrial settings, water hardness must be constantly monitored to avoid costly breakdowns in boilers, cooling towers, and other equipment that comes in contact with water. Hardness is controlled by the addition of chemicals and by large-scale softening with zeolite and ion exchange resins.

Softening

It is often desirable to soften hard water, as it does not readily form lather with soap. Soap is wasted when trying to form lather, and in the process, scum forms. Hard water may be treated to reduce the effects of scaling and to make it more suitable for laundry and bathing.

Process

A water softener, like a fabric softener, works on the principle of cation or ion exchange in which ions of the hardness minerals are exchanged for sodium or potassium ions, effectively reducing the concentration of hardness minerals to tolerable levels and thus making the water softer and giving it a smoother feeling
The most economical way to soften household water is with an ion exchange water softener. This unit uses sodium chloride (table salt) to recharge beads made of the ion exchange resins that exchange hardness mineral ions for sodium ions. Artificial or natural zeolites can also be used. As the hard water passes through and around the beads, the hardness mineral ions are preferentially absorbed, displacing the sodium ions. This process is called ion exchange. When the bead or sodium zeolite has a low concentration of sodium ions left, it is exhausted, and can no longer soften water. The resin is recharged by flushing (often back-flushing) with saltwater. The high excess concentration of sodium ions alter the equilibrium between the ions in solution and the ions held on the surface of the resin, resulting in replacement of the hardness mineral ions on the resin or zeolite with sodium ions. The resulting saltwater and mineral ion solution is then rinsed away, and the resin is ready to start the process all over again. This cycle can be repeated many times.
The discharge of brine water during this regeneration process has been banned in some jurisdictions due to concerns about the environmental impact of the discharged sodium.
Potassium chloride (softener salt substitute) may also be used to regenerate the resin beads. It exchanges the hardness ions for potassium. It also will exchange naturally occurring sodium for potassium resulting in sodium-free soft water.
Some softening processes in industry use the same method, but on a much larger scale. These methods create an enormous amount of salty water that is costly to treat and dispose of.
Temporary hardness, caused by hydrogen carbonate (or bicarbonate) ions, can be removed by boiling. For example, calcium bicarbonate, often present in temporary hard water, may be boiled in a kettle to remove the hardness. In the process, a scale forms on the inside of the kettle in a process known as "furring". This scale is composed of calcium carbonate.
Ca(HCO₃)₂ → CaCO₃ + CO₂ + H₂O
Hardness can also be reduced with a lime-soda ash treatment. This process, developed by Thomas Clark in 1841, involves the addition of slaked lime (calcium hydroxide — Ca(OH)₂) to a hard water supply to convert the hydrogen carbonate hardness to carbonate, which precipitates and can be removed by filtration:
Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃ + 2H₂O
The addition of sodium carbonate also permanently softens hard water containing calcium sulfate, as the calcium ions form calcium carbonate which precipitates out and sodium sulfate is formed which is soluble. The calcium carbonate that is formed sinks to the bottom. Sodium sulfate has no effect on the hardness of water.
Na₂CO₃ + CaSO₄ → Na₂SO₄ + CaCO₃

Alkalinity:

Alkalinity or A_T is a measure of the ability of a solution to neutralize acids to the equivalence point of carbonate or bicarbonate. Alkalinity is closely related to the acid neutralizing capacity (ANC) of a solution and ANC is often incorrectly used to refer to alkalinity. The alkalinity is equal to the stoichiometric sum of the bases in solution. In the natural environment carbonate alkalinity tends to make up most of the total alkalinity due to the common occurrence and dissolution of carbonate rocks and presence of carbon dioxide in the atmosphere. Other common natural components that can contribute to alkalinity include borate, hydroxide, phosphate, silicate, nitrate, dissolved ammonia, the conjugate bases of some organic acids and sulfide. Solutions produced in a laboratory may contain a virtually limitless number of bases that contribute to alkalinity. Alkalinity is usually given in the unit mEq/L (milliequivalent per liter). Commercially, as in the pool industry, alkalinity might also be given in the unit ppm or parts per million.
Alkalinity is sometimes incorrectly used interchangeably with basicity. For example, the pH of a solution can be lowered by the addition of CO₂. This will reduce the basicity; however, the alkalinity will remain unchanged.

References

1. ^ "The Measurement of pH - Definition, Standards and Procedures – Report of the Working Party on pH, IUPAC Provisional Recommendation"]. 2001. http://www.iupac.org/reports/provisional/abstract01/rondinini_prs.pdf.  A proposal to revise the current IUPAC 1985 and ISO 31-8 definition of pH.
2. ^ ^{a b c} Carlsberg Group Company History Page, http://www.carlsberggroup.com/Company/Research/Pages/pHValue.aspx
3. ^ University of Waterloo - The pH Scale, http://www.science.uwaterloo.ca/~cchieh/cact/c123/ph.html
4. ^ Nørby, Jens. 2000. The origin and the meaning of the little p in pH. Trends in the Biochemical Sciences 25:36-37., http://download.cell.com/trends/biochemical-sciences/pdf/PIIS0968000499015170.pdf
5. ^ Fundamentals of Analytical Toxicology, http://books.google.com.br/books?id=LBag6XlAJY0C
6. ^ Sørensen, http://www.geocities.com/bioelectrochemistry/sorensen.htm
7. ^ "pH". IUPAC Goldbook. http://goldbook.iupac.org/P04524.html. 
8. ^ Quantities and units – Part 8: Physical chemistry and molecular physics, Annex C (normative): pH. International Organization for Standardization, 1992.
9. ^ Definitions of pH scales, standard reference values, measurement of pH, and related terminology. Pure Appl. Chem. (1985), 57, pp 531–542.
10. ^ Nordstrom, DK et al (2000) Negative pH and extremely acidic mine waters from Iron Mountain California. Environ Sci Technol,34, 254-258.
11.  ^ DeZuane, John (1997). Handbook of Drinking Water Quality (2nd edition ed.). John Wiley and Sons. ISBN 0-471-28789-X. 
12. ^ ^{a b} C.M. Hogan, Marc Papineau et al. Development of a dynamic water quality simulation model for the Truckee River, Earth Metrics Inc., Environmental Protection Agency Technology Series, Washington D.C. (1987)
13. ^ USEPA. 1991. Guidance for water quality-based decisions: The TMDL process. EPA 440/4-91-001. U.S. Environmental Protection Agency, Office of Water, Washington, DC. ^ Boyd, Claude E. (1999). Water Quality: An Introduction. The Netherlands: Kluwer Academic Publishers Group. ISBN 0-7923-7853-9.
14. Holmes-Farley, Randy. "Chemistry and the Aquarium," Advanced Aquarist's Online Magazine. Alkalinity as it pertains to salt-water aquariums.
15. DOE (1994) ","Handbook of methods for the analysis of the various parameters of the carbon dioxide system in sea water. Version 2, A. G. Dickson & C. Goyet, eds. ORNL/CDIAC-74.