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  1. 1. EDUCATION HOLE PRESENTS ENGINEERING CHEMISTRY Unit-IV
  2. 2. Hardness of water ................................................................................................................... 2 Disadvantage of hard water .............................................................................................................3 Techniques for water softening ............................................................................................... 4 Calgon..............................................................................................................................................4 Ion Exchange or Zeolite ....................................................................................................................5 Other Softening Processes....................................................................................................................................5 Reverse-osmosis softening...............................................................................................................6 Ion-exchange resin devices...............................................................................................................6 Lime-Soda .............................................................................................................................................................7 Water treatment method for boiler feed by internal process ........................................................... 8 Phosphates-dispersants, polyphosphates-dispersants (softening chemicals)......................................................8 Natural and synthetic dispersants (Anti-scaling agents).......................................................................................8 Oxygen scavengers ...............................................................................................................................................8 Anti-foaming or anti-priming agents ....................................................................................................................9 Phase Rule ..................................................................................................................................... 10 Application to one component system................................................................................... 11 Hardness of water Hard water is water that has high mineral content (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 measures only 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. Iron in this case is important for, if present, it will be in its tervalent form, causing the calcification to be brownish (the color of rust) instead of white (the color of most of the other compounds).
  3. 3. Because it is the precise mixture of minerals dissolved in the water, together with the water's pH and temperature, that determines the behaviour of the hardness, a single-number scale does not adequately describe hardness. Descriptions of hardness correspond roughly with ranges of mineral concentrations: This scale is in substantial disagreement with the references. Very soft: 0-70 ppm 0-4 dGH Soft: 70-140 ppm 4-8 dGH Slightly hard: 140-210 ppm 8-12 dGH Moderately hard: 210-320 ppm 12-18 dGH Hard: 320-530 ppm 18-30 dGH Very hard: >530 ppm >30 dGH It is possible to measure the level of total hardness in water by obtaining a total hardness water testing kit. These kits measure the level of calcium and magnesium in the water. Temporary hardness test kits do not normally measure calcium and magnesium levels but normally use an approximation based on some form of alkalinity test. Measuring temporary hardness accurately would involve a series of tests to work out how much bicarbonates and carbonates are present and how much calcium and magnesium is present and what percentage combination there is. In most cases, the temporary hardness kit is a good approximation, but anions such as hydroxides, borates, phosphates can have quite an effect on temporary hardness test kits. Although most of the above measures define hardness in terms of concentrations of calcium in water, any combination of calcium and magnesium cations having the same total molarity as a pure calcium solution will yield the same degree of hardness. Consequently, hardness concentrations for naturally occurring waters (which will contain both Ca2+ and Mg2+ ions), are usually expressed as an equivalent concentration of pure calcium in solution. For example, water that contains 1.5 mmol/L of elemental calcium (Ca2+) and 1.0 mmol/L of magnesium (Mg2+) is equivalent in hardness to a 2.5 m. Disadvantage of hard water The disadvantages of having hard water is that if you have a boiler or a kettle and they are used for a long period of time then they will begin to clog up with fur and limescale. Another disadvantage is that when you have a bath with bubbles then the water will react with the soap so you won't get lather but a scum. Bathroom
  4. 4. Showerheads and spray-nozzles can become blocked; they can even clog the small holes completely and reduce their efficiency. The bathtub and sink seem to be the places where there is a visible soap scum build up. Without proper treatment his build up is very difficult to remove and may require a lot of cleaners and many applications. Scale can clog pipes and can decrease the life of toilet flushing units. Bathing Bathing with soap in hard water leaves a film of sticky soap curd on the skin. The film may prevent removal of soil and bacteria. Soap curd interferes with the return of skin to its normal, slightly acid condition, and may lead to irritation. Soap curd on hair may make it dull, lifeless and difficult to manage. Similarly, the insoluble salts that get left behind from using regular shampoo in hard water tend to leave hair rougher and harder to detangle. Laundering Clothes washed in hard water often look dingy and feel harsh and scratchy. The hardness minerals combine with some soils to form insoluble salts, making them difficult to remove. Soil on clothes can introduce even more hardness minerals into the wash water. Continuous laundering in hard water can damage fibers and shorten the life of clothes by up to 40 percent. Dishwashers When washing dishes, especially in a dishwasher, hard water may cause spotting and filming on your crockery. The minerals from hard water are released faster when it comes into contact with heat, causing an increase in the amount of spotting and filming that occurs. This problem is not a health risk, but it can be a nuisance to clean and reduce the quality of your crockery. Techniques for water softening Calgon Softening through chemical precipitation is similar to removal of turbidity by coagulation, flocculation, and sedimentation. There are many variations, but the typical process involves adding lime to raise the pH of water until it is high enough for reactions to occur which prompt hardness compounds to settle out of the water. The equipment used also resembles turbidity removal equipment - lime is added in the flash mixer, the water is flocculated, and then the hardness compounds precipitate out in the sedimentation basin. As mentioned above, groundwater is more likely to need softening than surface water is. Groundwater also may not need flocculation to remove turbidity, so the softening process can sometimes replace the turbidity removal process. If both turbidity removal and softening are required, then the two processes can occur simultaneously, using the same equipment.
  5. 5. Chemical precipitation using lime will remove carbonate hardness. If soda ash is added as well as lime, both carbonate and noncarbonate hardness may be removed. In either case, chemical precipitation does not remove all hardness from water. The hardness can be reduced as low as 30 to 40 mg/L using chemical precipitation, although the typical goal is 80 to 90 mg/L. We will discuss the chemical reactions which occur in lime-soda ash softening in a later section. Chemical precipitation is an effective softening process, but it does have some disadvantages. The process requires a lot of operator control to get an efficient result, which may make lime softening too operator-intensive for small treatment plants. The high pH used in lime softening can set colors in water and make them difficult to remove. Finally, lime softening produces large quantities of sludge which can create disposal problems. Ion Exchange or Zeolite Ion exchange softening, also known as zeolite softening, passes water through a filter containing resin granules. In the filter, known as a softener, calcium and magnesium in the water are exchanged for sodium from the resin granules. The resulting water has a hardness of 0 mg/L and must be mixed with hard water to prevent softness problems in the distributed water. Ion exchange softening does not require the flash mixer, flocculation basin, and sedimentation basin required for lime-soda ash softening. In addition, the process does not require as much operator time. Ion exchange softening is effective at removing both carbonate and noncarbonate hardness and is often used for waters high in noncarbonate hardness and with a total hardness less than 350 mg/L. However, ion exchange softening has its disadvantages as well. The calcium and magnesium in the hard water are replaced by sodium ions, which may cause problems for people with health problems who are not supposed to eat any salt. Softeners have to be backwashed in a manner similar to a filter, and the recharge water, known as brine, can cause disposal problems. Other Softening Processes Other processes can be used to soften water, but they are generally expensive and only used in rare circumstances. These alternative processes are listed below.
  6. 6. Reverse-osmosis softening Reverse-osmosis softening involves water being forced through a semi-permeable membrane. Calcium, magnesium, and dissolved solids are captured while the softened water is passed through the membrane. Ion-exchange resin devices Conventional water-softening appliances intended for household use depend on an ion-exchange resin in which "hardness ions" - mainly Ca2+ and Mg2+ - are exchanged for sodium ions. As described by NSF/ANSI Standard 44, ion exchange devices reduce the hardness by replacing magnesium and calcium (Mg2+ and Ca2+ ) with sodium or potassium ions (Na+ and K+ )."
  7. 7. Lime-Soda Chemical precipitation is one of the more common methods used to soften water. Chemicals normally used are lime (calcium hydroxide, Ca(OH)2) and soda ash (sodium carbonate, Na2CO3). Lime is used to remove chemicals that cause carbonate hardness. Soda ash is used to remove chemicals that cause non-carbonate hardness. When lime and soda ash are added, hardness-causing minerals form nearly insoluble precipitates. Calcium hardness is precipitated as calcium carbonate (CaCO3). Magnesium hardness is precipitated as magnesium hydroxide (Mg(OH)2). These precipitates are then removed by conventional processes of coagulation/flocculation, sedimentation, and filtration. Because precipitates are very slightly soluble, some hardness remains in the water--usually about 50 to 85 mg/l (as CaCO3). This hardness level is desirable to prevent corrosion problems associated with water being too soft and having little or no hardness. CO2 does not contribute to the hardness, but it reacts with the lime, and therefore uses up some lime before the lime can start removing the hardness. CO2 = carbon dioxide, Ca(OH)2 = calcium hydroxide or hydrated lime, CaCO3 = calcium carbonate, Ca(HCO3)2 = calcium bicarbonate, Mg(HCO3)2 = magnesium bicarbonate, MgCO3 = magnesium carbonate, Mg(OH)2 = magnesium hydroxide, MgSO4 = magnesium sulfate, CaSO4 = calcium sulfate, H20 - water. Na2CO3 = sodium carbonate or soda ash For each molecule of calcium bicarbonate hardness removed, one molecule of lime is used. For each molecule of magnesium bicarbonate hardness removed, two molecules of lime are used. For each molecule of non-carbonate calcium hardness removed, one molecule of soda ash is used. For each molecule of non-carbonate magnesium hardness removed one molecule of lime plus one molecule of soda ash is used.
  8. 8. Water treatment method for boiler feed by internal process Internal treatment can constitute the unique treatment when boilers operate at low or moderate pressure, when large amounts of condensed steam are used for feed water, or when good quality raw water is available. The purpose of an internal treatment is to • react with any feed-water hardness and prevent it from precipitating on the boiler metal as scale; • condition any suspended matter such as hardness sludge or iron oxide in the boiler and make it non-adherent to the boiler metal; • provide anti-foam protection to allow a reasonable concentration of dissolved and suspended solids in the boiler water without foam carry-over; • eliminate oxygen from the water and provide enough alkalinity to prevent boiler corrosion. In addition, as supplementary measures an internal treatment should prevent corrosion and scaling of the feed-water system and protect against corrosion in the steam condensate systems. During the conditioning process, which is an essential complement to the water treatment program, specific doses of conditioning products are added to the water. The commonly used products include: Phosphates-dispersants, polyphosphates-dispersants (softening chemicals) Reacting with the alkalinity of boiler water, these products neutralize the hardness of water by forming tricalcium phosphate, and insoluble compound that can be disposed and blow down on a continuous basis or periodically through the bottom of the boiler. Natural and synthetic dispersants (Anti-scaling agents) Increase the dispersive properties of the conditioning products. They can be: • Natural polymers: lignosulphonates, tannins • Synthetic polymers: polyacrilates, maleic acrylate copolymer, maleic styrene copolymer, polystyrene sulphonates etc. Oxygen scavengers Sodium sulphite, tannis, hydrazine, hydroquinone/progallol-based derivatives, hydroxylamine derivatives, hydroxylamine derivatives, ascorbic acid derivatives, etc These scavengers,
  9. 9. catalyzed or not, reduce the oxides and dissolved oxygen. Most also passivate metal surfaces. The choice of product and the dose required will depend on whether a deaerating heater is used. Anti-foaming or anti-priming agents Mixture of surface-active agents that modify the surface tension of a liquid, remove foam and prevent the carryover of fine water particles in the steam The softening chemicals used include soda ash, caustic and various types of sodium phosphates. These chemicals react with calcium and magnesium compounds in the feed water. Sodium silicate is used to react selectively with magnesium hardness. Calcium bicarbonate entering with the feed water is broken down at boiler temperatures or reacts with caustic soda to form calcium carbonate. Since calcium carbonate is relatively insoluble it tends to come out of solution. Sodium carbonate partially breaks down at high temperature to sodium hydroxide (caustic) and carbon dioxide. High temperatures in the boiler water reduce the solubility of calcium sulphate and tend to make it precipitate out directly on the boiler metal as scale. Consequently calcium sulphate must be reacted upon chemically to cause a precipitate to form in the water where it can be conditioned and removed by blow-down. Calcium sulphate is reacted on either by sodium carbonate, sodium phosphate or sodium silicate to form insoluble calcium carbonate, phosphate or silicate. Magnesium sulphate is reacted upon by caustic soda to form a precipitate of magnesium hydroxide. Some magnesium may react with silica to form magnesium silicate. Sodium sulphate is highly soluble and remains in solution unless the water is evaporated almost to dryness. There are two general approaches to conditioning sludge inside a boiler: by coagulation or dispersion. When the total amount of sludge is high (as the result of high feed-water hardness) it is better to coagulate the sludge to form large flocculent particles. This can be removed by blow-down. The coagulation can be obtained by careful adjustment of the amounts of alkalis, phosphates and organics used for treatment, based on the fee-water analysis. When the amount of sludge is not high (low feed water hardness) it is preferable to use a higher percentage of phosphates in the treatment. Phosphates form separated sludge particles. A higher percentage of organic sludge dispersants is used in the treatment to keep the sludge particles dispersed throughout the boiler water. The materials used for conditioning sludge include various organic materials of the tannin, lignin or alginate classes. It is important that these organics are selected and processed, so that they are both effective and stand stable at the boiler operating pressure. Certain synthetic organic materials are used as anti-foam agents. The chemicals used to scavenge oxygen include sodium sulphite and hydrazine. Various combinations of polyphosphates and organics are used for preventing scale and corrosion in feed-water systems. Volatile neutralizing amines and filming inhibitors are used for preventing condensate corrosion. Common internal chemical feeding methods include the use of chemical solution tanks and proportioning pumps or special ball briquette chemical feeders. In general, softening chemicals (phosphates, soda ash, caustic, etc.)
  10. 10. are added directly to the fee-water at a point near the entrance to the boiler drum. They may also be fed through a separate line discharging in the feed-water drum of the boiler. The chemicals should discharge in the fee-water section of the boiler so that reactions occur in the water before it enters the steam generating area. Softening chemicals may be added continuously or intermittently depending on feed-water hardiness and other factors. Chemicals added to react with dissolved oxygen (sulphate, hydrazine, etc.) and chemicals used to prevent scale and corrosion in the feed-water system (polyphosphates, organics, etc.) should be fed in the feed- water system as continuously as possible. Chemicals used to prevent condensate system corrosion may be fed directly to the steam or into the feed-water system, depending on the specific chemical used. Continuous feeding is preferred but intermittent application will suffice in some cases Phase Rule Minerals are the monitors of the physical and chemical conditions under which they formed. The occurrences of minerals, their parageneses (stable associations), types of reactions, and compositional variation (e.g. zoned minerals) all provide important information about geologic history and processes. Of particular importance to geologists are: • Estimates of pressure and temperature (geothermobarometry) • Estimates of other physico-chemical conditions such as acidity (pH) and oxidation state (eH) • Partial pressures of gases (e.g. fugacities of H2O, CO2 , etc.) • Partitioning of major and trace elements between phases (e.g. minerals, melts and/or fluids) to characterize and quantify petrogenetic processes; and • Use of minerals in geochronology and thermochronology Gibbs' Phase Rule provides the theoretical foundation, based in thermodynamics, for characterizing the chemical state of a (geologic) system, and predicting the equilibrium relations
  11. 11. of the phases (minerals, melts, liquids, vapors) present as a function of physical conditions such as pressure and temperature. Gibbs' Phase Rule also allows us to construct phase diagrams to represent and interpret phase equilibria in heterogeneous geologic systems. In the simplest understanding of phase diagrams, stable phase (mineral) assemblages are represented as "fields" (see colored areas on the figure to the right) in "P-T space", and the boundaries between stable phase assemblages are defined by lines (or curves) that represent reactions between the phase assemblages. The reaction curves actually represent the condition (or the locus of points in P-T space) where ΔGrxn =0; for more information on this point see Gibbs Free Energy. A solid understanding of Gibbs' Phase Rule is required to successfully master the applications of heterogeneous phase equilibria presented in this module. Gibbs Phase Rule is expressed by the simple formulation: P + F = C + 2, where P is the number of phases in the system A phase is any physically separable material in the system. Every unique mineral is a phase (including polymorphs); igneous melts, liquids (aqueous solutions), and vapor are also considered unique phases. It is possible to have two or more phases in the same state of matter (e.g. solid mineral assemblages, immiscible silicate and sulfide melts, immiscible liquids such as water and hydrocarbons, etc.) Phases may either be pure compounds or mixtures such as solid or aqueous solutions--but they must "behave" as a coherent substance with fixed chemical and physical properties. Application to one component system • The system is entirely composed of H2O, so there is only one component present. • The phases present represent three states of matter: liquid (water), solid (ice), and vapor (steam). All have distinct physical properties (e.g. density, structure--or lack of, etc.) and chemical properties (e.g. ΔGformation, molar volume etc.) so they must be considered distinct phases. • Note that there is only one point on this diagram where all three phases coexist in equilibrium--this "triple point" is also referred to as an invariant point; because P and T are uniquely specified, there are zero degrees of freedom. • Each of the curves represents a chemical reaction that describes a phase transformation: solid to liquid (melt/crystallization), liquid to vapor (boiling/condensation), solid to vapor (sublimation/deposition). There are three univariant curves around the invariant point; it is always the case that for a C-component system, there will always be C+2 univariant curves radiating around an invariant point. This relationship is further explained in the
  12. 12. unit on the Method of Schreinemakers. There is only one degree of freedom along each of the univariant curves: you can independently change either T or P, but to maintain two coexisting phases along the curve the second variable must change by a corresponding fixed amount. • There are three distinct areas where only ice, liquid, or vapor exit. These are divariant fields. T and P are both free to change within these fields and you will still have only one phase (a bit hotter or colder, or compressed or expanded, but nonetheless the same phase). • The end of the "boiling curve", separating the liquid to vapor transition, is called the "critical point". This is a particularly interesting part of the phase diagram because beyond this region the physico-chemical properties of water and steam converge to the point where they are identical. Thus, beyond the critical point, we refer to this single phase as a "supercritical fluid". Now, let's consider a simple 1 component system that describes the mineral phases in the aluminosilicate system: Phase diagram for the one component system Al2SiO5. • The entire system is defined by one component: Al2SiO5 (i.e. all the phases can be completely made of this one component) • There are three solid phases shown in this diagram: the polymorphs of Al2SiO5 andalusite, kyanite and sillimanite. • There is only one unique place on this diagram where all three phases can coexist in equilibrium--the invariant point at 3.8 Kb and 500o C; at this point there are zero degrees of freedom. • There are three univariant reactions on this diagram, each representing the phase transitions: andalusite = sillimanite, andalusite = kyanite, and kyanite = sillimanite. In each of these reactions, either pressure or temperature can be changed independently, but for the state of the system to remain the same (i.e. two solid phases coexisting in equilibrium), the other variable must change by a fixed amount to maintain the assemblage on the univariant curve -- so there is one degree of freedom. In a later section, we will see that the univariant curves represent the condition where ΔGrxn = 0 (i.e. the
  13. 13. intersection of the "free energy surface" with the Pressure-Temperature plane represented by the phase diagram). • There are three divariant fields in which only a single mineral phase is stable. Within these fields pressure and temperature may be changed independently without changing the state of the system--thus there are two degrees of freedom in the divariant fields.

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