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1. EDUCATION HOLE PRESENTS
ENGINEERING CHEMISTRY
Unit-IV

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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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.