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Environmental Ion Exchange
Principles and Design
Anthony M. Wachinski
Lime–Soda Softening
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Anthony M. Wachinski
Published online on: 29 Sep 2016
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2
Lime–Soda Softening
A fundamental knowledge of cold and hot lime–soda softening processes to understand and prepare bar charts is paramount in understanding the basic principles of
ion exchange.
HARDNESS
Hardness (any material that forms a precipitate with the sodium salt of fatty acids) is
classified as either permanent or temporary. The only chemical compound that forms
temporary hardness is calcium bicarbonate [Ca(HCO3)2], and it depends on the presence of dissolved carbon dioxide (CO2) in the water to exist. In the absence of CO2,
[Ca(HCO3)2] reverts to calcium carbonate (CaCO3). (Note: Calcium bicarbonate does
not exist in the dry form.) All other hardness is permanent.
Hardness can also be classified as carbonate or noncarbonate. Carbonate hardness
is classified as bicarbonate, carbonate, and hydroxide after the three forms of alkalinity. All other hardness is noncarbonate, for example, sulfate, nitrate, and chloride.
Powell (1954) defines hardness as “a phenomenon caused by the presence of the
multivalent cations, calcium and magnesium, in combination with bicarbonates, carbonates, sulfates, chlorides and nitrates.” In general, iron, aluminum, and manganese can also cause water to be hard, but these substances are not ordinarily present
in appreciable quantities. Carbonates are found only occasionally, in highly alkaline
waters. Nitrates are usually present in small concentrations and, on average, sulfates
likely exceed chlorides. There are, of course, many water supplies to which these
generalizations do not apply.
SOFTENING
What are the problems associated with the multivalent ions, calcium and magnesium, in water that warrant removal or substitution of these ions, that is, softening?
1. The loss of CO2 .by Ca(HCO3)2 forms CaCO3, which can scale pipes.
2. The reaction of hardness with soaps causes the soaps to be less effective. It
is sometimes cheaper to soften the water with lime and soda ash than to let
the soap do it.
3. The formation of precipitated solids in a boiler causes a reduction in the
heating efficiency of the boiler causing premature downtime and expense.
The removal of hardness was originally called zeolite softening; later the term softening, or water softening, was used. Common practice in the United States is to
identify water-softening procedures by the chemicals used in the process, thus the
term lime–soda softening. The processes are further classified as cold lime–soda
13
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14
Environmental Ion Exchange
softening, warm lime-soda softening, and hot lime–soda softening based on the
process operating temperature. Cold softening usually operates between 15°C and
60°C, warm softening between 60°C and 85°C, and hot softening between 90°C and
105°C. Cold softening was originally used for softening potable water. It was popular
until recently when municipalities stopped softening potable water. Warm and hot
lime softening are most often used in industrial water treatment. Hot lime softening
is most often used for treating make-up water for high pressure boilers that require
very low hardness feed waters. Hot lime softening and warm lime softening are
often used in oil extraction facilities to soften produced water to prepare boiler feed
water for once through steam generators (OTSGs). Both practically generate water
with almost the same level of hardness. In theory, hot lime softening produces water
with lower residual hardness, requires less to no coagulant, removes dissolved gases,
and does a better job in reducing silica than cold softening. Lime-soda softening is
sometimes used to remove those constituents that negatively impact the recovery of
diffusive membrane processes, nanofiltration, and reverse osmosis.
This chapter addresses the principles common to cold, warm, and hot lime–soda
softening and provides a detailed discussion on the preparation and use of bar charts.
BASIC PRINCIPLES
Water softening by chemical addition uses classical precipitation methods, that
is, rapid mixing, which converts hardness to a low solubility salt, flocculation, to
agglomerate the particles into larger flocs, sedimentation to settle the particles and
flocs. Recarbonation to add alkalinity to the finished water to adjust the Langelier
saturation index, and filtration (if required), to reduce particulate turbidity and associated hardness.
The process converts the soluble multivalent cations responsible for hardness into
insoluble compounds that can be reduced to about 35 mg/L as CaCO3 because of the
solubilility of CaCO3 and Mg(OH)2 by traditional solid separation techniques without excessive use of chemicals. The minimum calcium hardness that can be achieved
is about 30 mg/L as CaCO3, and magnesium hardness is about 10 mg/L as CaCO3.
For drinking water, a final total hardness on the order of 75–120 mg/L as CaCO3 is
targeted. In practice, residual hardness ranges from 50 to 100 mg/L for well-run
plants and double that concentration for plants not well run.
Lime Ca(OH)2 and soda ash (Na2CO3) can be added to the water to obtain maximum precipitation of Ca++ and Mg++. Chemical softening or chemical precipitation
softening has been the preferred method because it is the cheapest. Whether it will
remain a viable process is yet to be decided. The sludge produced by this process
has little value and is difficult to dewater and dispose of, resulting in high trucking
fees and disposal costs. As a result, the process is often replaced by nanofiltration.
Nanofiltration is a membrane process often used to remove divalent salts from
water, that is, membrane softening. Nanofiltration membrane systems, that is, membrane softeners retain dissolved organic compounds in the range of 200–400 Daltons,
essentially all multivalent cations and anions, and a fraction of the monovalent species. They operate in the 100–200 pounds per square inch (psi) (6.9–13.8 bar) range.
These membranes achieve greater than 95% rejection of calcium and magnesium.
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15
Lime–Soda Softening
Membrane softeners are becoming popular as industrial replacement technology for
lime-soda softening in house and as point of use (POU) and point of entry (POE)
softening technology.
The first reaction is an unwanted reaction with dissolved CO2.
CO2 + Ca(OH)2 → CaCO3 + H2O
(2.1)
Although a totally useless reaction, it cannot be prevented. Because of this reaction,
CO2 must be added to the water after softening. This process is called recarbonation.
The next reactions are
Ca(HCO3)2 + Ca(OH)2 → 2CaCO3 + 2H2O
(2.2)
Mg(HCO3)2 + 2Ca(OH)2 → Mg(OH)2 + 2CaCO3 + 2H2O
(2.3)
A method is required to get the lime–soda reactions to move to completion. Both
Ca++ and Mg++ carbonates are soluble to some degree, and thus a driving force is
needed to force completion of the reaction. The usual method is to add excess chemical thus shifting the equilibrium toward the carbonate precipitates. A pH value of
10.3 is the minimum solubility for Mg(OH)2, and a pH of 9.3 is the minimum solubility for calcium (this minimum is 8–10 mg/L). Thus, to get both to a minimum,
a chemical “trick” must be used.
Adding excess lime and excess soda ash to the reaction can lower the concentrations of Ca++ and Mg++ to about 8 mg/L or a total hardness of 16 mg/L as a minimum. Note that the hot process allows lower levels of hardness because at higher
temperatures CO2 is less soluble and thus does not load the reactions in favor of the
soluble bicarbonates.
The lime–soda ash process is used in produced water treatment and in boiler feed
operations to reduce the silica concentration. When sodium aluminate and ferric
chloride are added, the precipitate will include CaCO3 and a complex with silicic
acid, aluminum oxide, and iron.
With the hot lime silicic acid removal process at 60°C–70°C, silica can be reduced
to 1 mg/L by adding a mixture of lime and porous magnesium oxide.
With lime softening, barium, strontium, and organic substances are also reduced
significantly. The process requires a reactor with a high concentration of precipitated
particles serving as crystallization nuclei.
After lime–soda softening, the water is supersaturated with CaCO3, which will
precipitate or coat piping after leaving the process. Thus the water must be stabilized
before discharging to the distribution system. The usual method for stabilization is
recarbonation. In recarbonation, CO2 is added to the water to lower the pH, forming
bicarbonates from the carbonates and ensuring that the materials will precipitate.
All of the hydroxide and carbonate alkalinity that could combine with calcium or
magnesium must be converted to bicarbonate to be sure that there will not be precipitation later in the distribution lines.
Care must be taken in the recarbonation process to ensure that too much CO2 is
not introduced into the water. The required amount of CO2 was determined from
a bar diagram by calculating the amount required to destroy only the hydroxide
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16
Environmental Ion Exchange
alkalinity. Today software programs are used. Conversion of Na2CO3 to NaHCO3
is possible but not necessary due to the high solubility of Na2CO3. If all of the carbonate is converted to bicarbonate, there is an excellent chance of overcarbonation
with a resultant drop in pH to below 8.4. With the pH below 8.4 and excess CO2 in
the water, the water will become aggressive and cause a corrosion problem in the
distribution system.
The usual method of recarbonation in all but the smallest plants is to burn methane to produce CO2 and waste the heat. Small plants typically use a controlled feed
of sulfuric acid to convert some of the CaCO3 to CaSO4, which has a higher solubility, and forms H2CO3. (Note: The same reaction applies to MgCO3.)
The only way to reduce the total dissolved solids (TDS) of a water by precipitation is if the TDS, or part of it, is in the form of bicarbonates. The bicarbonates can
be precipitated as insoluble carbonates, whereas sulfates and chlorides cannot be
precipitated at any reasonable concentration.
BAR CHARTS
Bar charts are used to make regenerant calculations and to “check” laboratory analyses. An understanding of bar charts makes the more difficult concepts associated
with ion exchange easier to grasp.
An order of preference exists in nature for combining anions and cations. For
cations, the order is as follows:
Aluminum
Calcium
Hydrogen
Iron
Magnesium
Potassium
Sodium
For anions, the order is as follows:
Bicarbonate
Carbonate
Chloride
Fluoride
Hydroxide
Nitrate
Sulfate
To plot a bar chart, first express all of the constituents in similar units (e.g., mg/L as
CaCO3 or Eq/L). Two bars are constructed equal to the scale length of total anions
or cations. Plot cations on the top bar and anions on the bottom bar as shown above.
(Note: In most waters, some CO2 is dissolved in the sample but is usually not reported
in the analysis. The quickest and most accurate way to obtain the CO2 concentration
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17
Lime–Soda Softening
is by using the TDS and pH with the chart in any recent edition of Standard Methods
for the Examination of Water and Wastewater. Standard Methods for the Examination
of Water and Wastewater is comprehensive reference that covers all aspects of water
and wastewater analysis techniques. Standard Methods is a joint publication of the
American Public Health Association (APHA), the American Water Works Association
(AWWA), and the Water Environment Federation (WEF).)
The following example illustrates how to construct a bar chart. Because both
phenolphthalein (P) and methyl orange (MO) alkalinity are given, we know that the
pH is greater than 8.3. (See Chapter 1 for a review of alkalinity relationships.) The
carbonate (CO3−−) alkalinity is 140 mg/L expressed as CaCO3 and the bicarbonate
alkalinity is 235 mg/L expressed as CaCO3.
Example
Bar Chart Problem
Given:
P alk =
MO alk =
Ca++ =
Mg++ =
Na+ =
K+ =
SO=4 =
Cl− =
70 mg/L as CaCO3
375 mg/L as CaCO3
80 mg/L as Ca++
24 mg/L as Mg++
30 mg/L as Na+
27 mg/L as K+
14.4 mg/L as SO4=
7.1 mg/L as Cl−
Ca
Mg
Na
K
SO4
Cl
CO3
Atomic Weight
Equivalent Atomic
Weight
40
24
23
39
96
35.5
60
20
12
23
39
48
35.5
30
Find: The hypothetical combinations of compounds and prepare a bar chart.
Solution
Because
=
P < 1 2 MO, CO=
2=
P (2)(70) = 140 mg /L as CaCO3
3
HCO3− = T − 2P = 375 − (2)(70) = 375 − 140 = 235
= 235 mg/L as CaCO3
where T is the total alkalinity or methyl orange alkalinity.
Express constituents as calcium carbonate (CaCO3)
Ca = 80 ×
50
= 200 mg/L as CaCO3
20
Mg = 24 ×
50
= 100 mg/L as CaCO3
12
Na = 30 ×
50
= 65 mg/L as CaCO3
23
K = 27 ×
50
= 35 mg/L as CaCO3
39
Environmental Ion Exchange
0
200
300
Ca++
Mg++
HCO3−
FIGURE 2.1
Hydroxide
Bicarbonate
Carbonate
Sulfate
Chloride
Fluoride
Nitrate
Na+
CO3=
235
Using order of preference:
Iron
Aluminum
Calcium
Magnesium
Sodium
Potassium
Hydrogen
365
React in
this order
400
K+
SO4=
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18
Cl−
375 390 400
Based on rationale of solubilities
Example problem final bar chart.
SO=4 = 14.4 ×
Cl− = 7.1×
50
=15 mg/L as CaCO3
48
50
= 10 mg/L as CaCO3
35.5
Figure 2.1 shows the final bar chart.
PROCESS EQUATIONS
The following equations summarize the lime–soda softening computations. Before
softening can occur, lime will be consumed by any free carbon dioxide present in
the water:
2CO2 + Ca(OH)2 → Ca(HCO3 )2
Carbon dioxide + Calcium hydroxide
→ Calcium bicarbonate ( hydrated lime)
(2.4)
This first reaction produces calcium bicarbonate, which adds hardness.
The most important reaction in water softening is the conversion of soluble calcium bicarbonate to the relatively insoluble carbonate, which can be removed by
settling and filtration. Sufficient lime is added to react with the Ca(HCO3)2, which
precipitates the carbonate according to the following reaction:
Ca(HCO3 )2 + Ca(OH)2 → 2CaCO3 ↓ +2H 2O
Calcium bicarbonate (hydrated lime) + Calcium hydroxide
(2.5)
→ Calcium carbonate + ( Water)
Partial softening may be accomplished where complete softening is not justified by
adding enough lime to convert only part of the calcium bicarbonate to carbonate.
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Lime–Soda Softening
19
Of the calcium carbonate thus formed, some will remain in solution (according to its
solubility), and the remainder will precipitate.
If magnesium or sodium bicarbonate is present, lime consumption for complete
treatment is shown by the following family of reactions:
Mg(HCO3 )2 + Ca(OH)2 → MgCO3 ↓ + CaCO3 ↓ + 2H 2OS
Magnesium bicarbonate (hydrated lime) + Calcium hydroxide
(2.6)
→ Magnesium carbonate + Calcium carbonate + ( Water )
2NaHCO3 + Ca(OH)2 → Na 2CO3 + CaCO3 ↓ +2H 2O
Sodium bicarbonate (hydrated lime) + Calcium hydroxide
(2.7)
→ Sodium carbonate + Calcium carbonate + ( Water )
Sodium and magnesium carbonates formed by the foregoing reaction are soluble and
are not precipitated. Therefore, the additional lime used does not directly remove
an equivalent amount of hardness. However, the soluble carbonates that are formed
improve the performance of the process because of the excess of soluble carbonates
produced.
Lime precipitates magnesium hydroxide:
MgCO3 + Ca(OH)2 → Mg(OH)2 ↓ + CaCO3 ↓
(2.8)
Magnesium carbonate + Lime → Magnesium hydroxide + Calcium carbonate
MgSO4 + Ca(OH)2 → Mg(OH)2 ↓ + CaSO 4
(2.9)
Magnesium sulfate + Lime → Magnesium hydroxide + Calcium sulfate
MgCl2 + Ca(OH)2 → Mg(OH)2 ↓ + CaCl 2
(2.10)
Magnesium chloride + Lime → Magnesium hydroxide + Calcium chloride
Soluble magnesium carbonate uses one equivalent of lime in its formation from
magnesium bicarbonate (see Equation 2.5) and requires a second equivalent for its
conversion to the hydroxide. Magnesium sulfate and magnesium chloride, when precipitated with lime, produce a corresponding calcium salt, which adds to the calcium
noncarbonate hardness of the water (Equations 2.8 and 2.9).
In general, magnesium hardness can be left in solution or removed by adding
more lime, according to the requirements for the finished water. However, when lime
is fed for the purpose of removing calcium bicarbonate hardness, some magnesium
precipitates, and, as it consumes equivalent lime, the dosage must be increased to
meet this demand. Such “incidental” magnesium removal may vary from 10% of
the initial magnesium to much higher amounts where the total lime dosage is heavy.
The fourth fundamental family of reactions in lime–soda softening is the conversion by soda ash of soluble calcium sulfate and calcium chloride to insoluble calcium
carbonate.
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Environmental Ion Exchange
CaSO 4 + Na 2CO3 → CaCO3 ↓ + Na 2SO 4
Calcium sulfate + Soda ash → Calcium carbonate + Sodium sulfate
(2.11)
CaCl 2 + Na 2CO3 → CaCO3 ↓ + 2NaCl
(2.12)
Calcium chloride + Soda ash → Calcium carbonate + Sodium chloride
The calcium carbonate thus formed precipitates along with that produced from
calcium bicarbonate.
Soda ash is added not only to reduce the calcium noncarbonate hardness, but
also to remove magnesium noncarbonate hardness. Magnesium sulfate and chloride,
upon addition of lime, form an equivalent amount of calcium noncarbonate hardness,
which must then be precipitated with soda ash. In other words, the soda ash dosage
is based on the total amount of noncarbonate hardness, and no distinction need be
made between calcium and magnesium salts in estimating treatment requirements.
Four fundamental processes are thus involved in softening with lime and soda
ash, resulting in the formation of calcium carbonate and magnesium hydroxide. The
final hardness of the water will be governed by the solubility of these compounds
under the existing temperature conditions and other substances present in the finished or treated water.
Example
Lime–Soda Softening Problem
Given:
pH = 7.1
CO2 = 11 mg/L
Ca++ = 80 mg/L as Ca
Mg++ = 48 mg/L as Mg
Alk = 340 mg/L as CaCO3
SO4 = 106 mg/L as SO4=
Cl− = 36 mg/L as Cl−
Find: Treat 4 million gallons per day (MGD) to a residual hardness of 12 mg/L of
Ca as CaCO3 and 12 mg/L of Mg as CaCO3.
Solution
1. In terms of CaCO3:
Ca++ = 80 × 50/20 = 200 mg/L
Mg++ = 48 × 50/12 = 200 mg/L
Na+ = 46 × 50/23 = 100 mg/L
HCO3– = 340 mg/L
SO4= = 106 × 50/48 = 110 mg/L
Cl− = 36 × 50/35.5 = 50 mg/L
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21
Lime–Soda Softening
2. Bar graph
200
Ca
500
400
Mg
Na
HCO3
SO4
340
Cl
450
Species:
Ca(HCO3)2 = 200 mg/L as CaCO3
Mg(HCO3)2 = 140 mg/L as CaCO3
MgSO4 = 60 mg/L as CaCO3
Na2SO4 = 50 mg/L as CaCO3
NaCl = 50 mg/L as CaCO3
Chemicals:
1. CO2 + Ca(OH)2 → CaCO3 + H2O
Set-up proportions:
44 74 44 100
= =
S1
11 L1 11
L1
=
(11)(74)
(100)(11)
S1
=
44
44
L1 = 18.5mg/L
S1 = 25mg/L
2. Ca(HCO3)2 + Ca(OH)2 → 2CaCO3 + 2H2O
100 74 2 × 100
200 L2
S2
L2 = 148 mg/L
S2 = 400 mg/L
In this problem, molecular weight of Ca(OH)2 and CaCO3 are given
as Ca(OH)2 and CaCO3. The concentration of Ca(HCO3)2 is given as
CaCO3. So we have to change the 200 mg/L as CaCO3 to be expressed
as Ca(HCO3)2, that is, multiply by EW / 50. This part of the proportion
would be
162 (molwt) 74
=
200 × (81/50) L2
but 162 × (50 / 81) = 100 ⇒ will always be 100.
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Environmental Ion Exchange
3. Mg(HCO3)2 + 2Ca(OH)2 → 2CaCO3 + Mg(OH)2 + 2H2O
100 2 × 74 2 × 100 58
140 L3
S3
S4
L3 = 207 mg/L
S3 = 280 mg/L S4 = 81mg/L
4. MgSO4 + Ca(OH)2 → Mg(OH)2 + CaSO4
100 74 58
60 L4 S5
=
L4 44
=
mg/L S5 35mg/L
5. CaSO4 + Na2CO3 → CaCO3 + Na2SO4
100 106
=
60
A1
100
A6
=
=
A1 64
mg/L S6 60 mg/L
A6 = 60 mg/L
To decrease Mg++ from 33 to 12 mg/L as CaCO3
62mg/L of OH− alk as CaCO3 required
−12mg/L of OH− alk as CaCO3 to remain
50mg/L of OH− alk must be added
We will add the OH− as Ca(OH)2.
50mg/L of OH− alk as CaCO3 ×
37
(equivalent wt. Ca(OH)2 )
50
= 37 mg/L of excess lime must be added
Expressed as lime, L5 = 37 mg/L
But we have added Ca++ and must remove it!
Consider changing Ca++ to CaCO3
Ca(OH)2 + Na2CO3 → CaCO3 + 2NaOH
Set-up proportion:
74 106
=
37
A2
74 A2 =(37)(106)
A2 =
(37)(53)
37
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Lime–Soda Softening
A 2 = 53 mg/L of Na2CO3 must be added.
To lower the concentration of Ca++ from 35 to 12 mg/L
59 mg/L of CO3++ alkalinity as CaCO3 is required.
12 mg/L of CO3++ alkalinity as CaCO3 remains
47 mg/L of CO3++ as CaCO3 must be added
To do this, we will add the carbonate alkalinity using sodium carbonate (Na2CO3).
47 mg/L of Na2CO3 as CaCO3 must be converted to Na2CO3 expressed as
Na2CO3.
47mg/L ×
53
= 49.8mg/L
50
use50mg/L
A3 = 50mg/L
Lime required = L1 + L2 + L3 + L4 + L5
18.5 + 148 + 207 + 44 + 37 = 454.5mg/L = 15,162lb/day of 100% lime
Chemicals:
Soda ash required = 64 + 53 + 50 = 167mg/L = 5571lb/day of 100% soda ash
To determine the pounds of sludge generated, add “S”s
Sludge is S1 + S2 + S3 +S4 + S5 + S6
Sludge is 25 + 400 + 280 + 81 + 35 + 60 = 881 mg/L = 29, 390 pounds of
dry solids. At 3% solids = 979,000 pounds, 490 tons per day of 3% solid sludge.
BIBLIOGRAPHY
James M. Montgomery, Consulting Engineers, Inc., Water Treatment Principles and Design,
John Wiley and Sons, New York, pp. 291–299, 1985.
Metcalf and Eddy, Wastewater Engineering, Treatment, and Reuse, 4th ed., McGraw-Hill,
New York, 2004.
Powell, S. T., Water Conditioning for Industry, McGraw-Hill, New York, 1954.
Standard Methods for the Examination of Water and Wastewater, 18th ed., APHA, AWWA
and WEF, 1992.
Wachinski, A. M., Ion Exchange Treatment for Water, American Water Works Association,
Denver, CO, November 2004, Preface.
Wachinski, A. M., Membrane Processes for Water Reuse, McGraw Hill, New York, 2013,
Chapter 3.
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