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A Study of the Mechanism of

Water Softening with Lime

A Thesis Submitted to

The Faculty of
MICHIGAN STATE COLLEGE
of

AGRICULTURE AND APPLIED SCIENCE

By

I \n
._ ‘ T“)
I' l’ ' ‘1?"

:J' - A.
01 yr"
Alfred M.F.Johnson

Candidate for the Degree of

Bachelor of Science

March 1935

Acknowledgement

 

At this time I want to thank Mr. Edward F. Eldridge for
his very helpful criticisms and suggestions Which he so
kindly made from time to time during the course of this
investigation.' Also for his help in securing the sample
of river water from Midland, Michigan Which was used in this

work.

 

I am.also indebted to the Civil Engineering Department
L of Michigan State College for the use of the laboratory and

the chemicals and equipment needed for the analysis.

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TABLE OF CONTENTS

Introduction . . . . . . . . . . . . . . . . Page 1
Discussion of the Problem . . . . . . . . . . " ll
Tabulated Data and Results . . . . . . . . . . " 20
Graphs . . . . . . . . . . . . . . . . . . . . " 22
Discussion of Graphs . . . . . . . . . . . . . “ 23

Conclusion . . . . . . . . . . . . . . . . . . " 36

Bibliography . . . ... . . . . . . . . . . . .

A STUDY OF THE
MECHANISM OF WATER SOFTENING WITH LIME

Hardness in water is defined as that quality of the
water which produces a curd of insoluble material when used
with soap and which produces a scale when used in heating
and power equipment. It is due to the presence of salts of
calcium and magnesium.

It has long been recognized that all water supplies are
not of the same quantity of hardness. Some waters are that
is called hard and others are known as soft waters. This is
only one of many general classifications of water. A further
classification, and one Which gives a more comprehensive idea
of the waters in question, is the grouping of them by the
sources of supply, of which the principal ones are as follows:
SOFT WATERS

Rain Water-

Rain water is seldom pure because in its descent through
the atmosphere to the earth it takes up soluble gases, car-
bonic acid, dirt, dust, etc. In manufacturing districts, es-
pecially near smelters and chemical plants, the quantity of
such materials may be large.

Snow Wat er-

Snow water is practically the same as rain water, the
chief difference being in the temperature at the time of
precipitation.

Surface Water-

Surface water is rain water and water condensed from

(2)

the atmosphere which has run over the surface of the ground
and taken up impurities. It may contain some ground water.
River Water-

River water is surface water which has run off from.the
earth to its natural channel on its way to the ocean. It
usually contains ground water. Usually it is contaminated
by organic matter such as leaves, coloring matter, etc., as
well as soluble mineral matter. The quantity of these mat-
erials carried by river water depends upon the location of
the stream and its source. Many rivers are used for the dis-
posal of sewage and industrial wastes and may become highly
contaminated.

Lake Water-

Lake water is practically the same as river water. It
may contain a greater percentage of ground water than the
rivers if it be from what is called "a spring-fed lake."
Lakes, especially the smaller inland lakes, usually do not
receive as much sewage or industrial wastes because there is
less movement of the water in a lake than in a river and the
use of the lake as a means of disposal would soon create a
nuisance, both from.the aesthetic and the sanitary viewpoint.

The above water supplies are generally considered as being
soft water supplies except where ground water may have become
mixed with surface, river, or lake water and caused the resul-

tant water to become hard.

(3)

Hard‘Waters

 

Ground‘Water-

Ground water is water which comes out of the earth.

It often contains gaseous materials (principally carbon di-
oxide) given off by decaying vegetable matter. The carbonic
acid, which is produced when carbon dioxide dissolves in
water, dissolves the limestone in the earth and causes the
water to become hard.

Ground water may come from artesian wells and springs
or may be pumped from the ground. The source of this water
may be either in the surface subsoil or in the underlying
rock. Ground waters in general are much harder than surface

waters.

Hardness in water is divided into two classes--temporary

and permanent;

‘Tgmpgrary Hardness

Temporary hardness is that hardness which is removable
by boiling the water. It is caused by the presence of bicar-
bonate of calcium, Upon boiling the calcium is precipitated
from solution in the form of calcium carbonate.
Permanent Hardness

Permanent hardness is that hardness which is not re-
movable by boiling. It is due chiefly to the presence of

sulfates and chlorides of calcium and magnesium.

In order that these hardness-causing compounds and their

(4)

actions may be better understood, a brie{?discussion of each
of the principal ones follows:

Calcium carbonate, (0a 003), commonly known as limestone,
marble, or chalk, does not form a hard scale in boilers and
heating equipment but causes much of the soft mud found in
the bottoms of these units. In combination with other sub-
stances, however, it may form a very hard scale. 'It is read-
ily soluble in carbonic acid or in water containing carbonic
acid gas, (002). Carbonic acid is driven off when the temp

perature of the water reaches about 200 F. and the calcium

carbonate is precipitated.
Magnesium carbonate, (Mg 003), is a toilet preparation,

its common form being more soluble in.water than chlcium.car-

bonate. It is held in solution be carbonic acid and may be
precipitated by the use of slaked lime. It precipitates out

in boilers as magnesium hydrate and magnesium oxide. The
scale formed is a very good non-conductor and is used as an
insulating covering for boilers and pipes. This shows the
scale to be exceptionally undesirable.

Calcium sulfate (gypsum) forms the hardest form of scale
of these encountered in water supplies. It sticks tightly

to the walls of the boilers and is also a very poor conductor

of heat, causing a tremendous waste of heat in boiler opera-

tion.

Magnesium sulfate, (Mg 394), or epsom salt, is very

easily soluble in cold or warm water. It does not form

scale in boilers but when broken.up by lime salts, it forms

(5)

a scale composed of magnesium and calcium.aulfate. It has a
very corrosive action on boiler iron and fittings.

Calcium chloride (CaClg) is very soluble in natural water.
It is corrosive but does not form a scale unless in combina-
tion with MgSO4 when it causes a hard scale to form due to a
transfer of acids and in addition forms magnesium chloride
which is very corrosive.

Magnesium.chloride is very corrosive in boiler water,
causing pitting and grooving. It is very soluble in water
and is the destructive element in sea water. It is usually
found to accompany sodium chloride when the latter is found
in a natural water supply.

Carbon dioxide (002) is a gas found in all natural waters.
It forms carbonic acid with water which increases the solu-
bility of calcium and magnesium carbonates and is also favor-
able to pitting and corrosion. It can be completely removed.

The above described compounds are the principal offenders
in causing water to be hard. There are other objectionable
compounds, as of iron, aluminum, etc. present in water, but
they are usually present in such small quantities as to be
insignificant from the standpoint of hardness. They are
usually objectionable because of coloring the water or be-
cause of the staining of fixtures. The process of removing
them will not be considered here except to mention that it
is possible and is practiced in some cases.

While the chief objection to hard water arises from the

fact that it causes an excessive amount of soap to be used in

(6)

the ordinary processes of washing in the toilet, kitchen and
laundry, there are several conditions in industry where soft,
pure water is not only a matter of economy but an actual
necessity. A few of the more important of these cases will
be mentioned here.

Water for Concrete and Construction Work-

This water, while not required to be soft, should be

free from all oil, acids, strong alkalies and vegetable or
organic matter. Under certain conditions and within limits

Cs/JHE/e
the presence of sodium or calcium.may be favored as an aid

A.
to the curing of concrete in cold weather. The use of such
a water is a matter of debate among engineers and in case of
its use, it should be thoroughly analyzed before using.
Water in Brewing-

Brewing is a revived-industry in this country. Soft
water aids in.hte preparation of malt and adds to the clear-
ness of beer. For brewing purposes the water should be
colorless, odorless, insipid and free from organic material.
Mineral salts, in general, are harmful although some help
the quality of beer. Sulfates and a very small percentage
of salt are generally to be preferred in the water used.

Too much salt or carbonates of calcium or magnesium give a
bitter taste to beers and ales.
Water in Dyeing and Bleaching-

These industries are especially particular about the

quality of the water they use. Certain dyes and dyestuffs

require that the water used be of a certain quality and this

('7)

same water may be absolutely worthless for use with some
other dyes and may even cause a complete change of color of
the dye. In general, hard water has a tendency with aliza-
rine, direct and basic, dyes to throw down the color itself
as an insoluble precipitate and when iron is present, it
causes red stains. Hard water used to wash off the loose
dyes, and to which no corrective can be applied also has an
injurious effect on the fiber and colors of the cloth being
treated.

Water in the Tanning Industry-

Quicklime and water are first used to get the roots of
the hair out of the leather hide. Water containing calcium
carbonate is not good for this purpose, as it prevents the
absorption of the tannin, and may also cause brownish stains.
Magnesium.carbonate acts in a very similar manner. Calcium
and magnesium sulfates are very desirable, and in fact, are
necessary to the tanning water. Carbonic acid is also very
advantageous.

Water in Ice-making-

Ice to be used for cooling and which will eventually
be taken internally should be pure bacteriologically. This
type of purification is again outside the scOpe of this work
and will only be mentioned at this time. However, it is de-
sirable to have the water used for ice-making colorless and
odorless. It is also usually desirable that it contain a
moderate amount of air and mdneral matter.

Water in Paper-making-

(8)

Paper making requires a large amount of high quality
water. Fine papers require a water free from suspended or
dissolved matter. The Chinese have for centuries prepared
their water for paper making by adding alum to coagulate the
suspended and colloidal matter. The use of waters containing
iron salts results in a brown coloring, and lime and magne-
sium compounds interfere with the process of sizing the
paper. This effect is more noticeable with the vegetable
size now used than with the animal size formerly in use.

Water for Heating and Power Equipment-

The effect of hard waters on boilers has been discussed
before in this work. It should be noted that in the case of
railroad locomotives with clean boilers when using hard water
required calking at the end of the third week and also needed
thirty-three calkings in seven months, at which time the
flues were worn out. After changing to soft water the same
locomotives required only six calkings in fifteen and one-
half months.

Water in Sugar Refining-

Sugar syrups are filtered through animal charcoal and
if calcareous water is used to wash the charcoal filter the
charcoal soon loses its good qualities. Magnesium.chloride
is especially bad in sugar making as hydrochloric acid is
liberated which would soon cause trouble.

From.the preceding discussion it is evident that "soft”
water, either natural, as rain water, or a softened water is

essential for best and most economical use in industry and

(9)

the home.

The supply of natural rain water is quite obviously lim»
ited and of both uncertain quality and quantity. Hence,
practically all the water used must be a treated, or soft-
ened water.

It must be remembered that a hardness of zero p.p.m.
(parts per million) of hardness is not as desirable as the
foregoing discussion would seem to indicate. Water of zero
p.p.m. of hardness is undesirable in many cases because it
makes it difficult to rinse soap from the hands or from
clothes being laundered. Neither is it economically prac-
tical to soften the water supply for a city to zero p.p.m.
of hardness, due to the fact that it is quite difficult to
remove the last few degrees of hardness. Water having from
50-75 p.p.m. of hardness is usually considered as being soft
enough for a public water supply. If the water is of from
75-150 p.p.m. of hardness it would be considered moderately
hard but would still not be unduly hard or interfere with
its use for most purposes or arouse a demand for softening.
Hardness of more than 150 p.p.m. is noticed by nearly every-
one and will likely arouse a public demand for Softening.

Hunicipal softening may be done by either of two common
methods ;

(1) One is zeolite sand softening, in Which an exchange
of the sodium in the sand and the calcium.and magnesium.in
the water takes place. This method is capable of removing

—all the calcium and magnesium hardness and it also removes

(10)

the iron. It is satisfactory and economical to use for
small to mediumpsized water supplies.

(2) The other is lime, or lime and soda, softening, which
is better adapted to large supplies than is zeolite and hence
is more often used for municipal treatment of water supplies.

There is a method of softening by removal of the temper-
ary hardness by heat but it is used only in special cases
and, for obvious reasons, only on small supplies. It is used

in some boiler installations and under similar circumstances.

(11)

DISCUSSION OF THE PROBLEM

For my problem I chose to make a study of the softening
of water by the use of lime (as Ca (0H)2) alone, without the
addition of soda ash (sodium carbonate) and to observe the
results obtained by the different percentages of the theoret-
ical 100 % softening by the lime. In other words, to deter-
mine the mechanism of the lime softening reaction.

The equations expressing the reactions occurring when lime
is added to remove calcium and magnesium from water are as
follows:

(1) Ca (3003)2 + 03(0H)2 =,2 Ca 0034);.2 H20
(2) M8 (H003)2 4 Ca (OH)2 : mgcoa'+ CaCOSJ *.2 320

(5) H3003 +Ca(0H)2

Ma(on)2* 4 Cases

(4) 113304 + Ca(CH)2

Mg(0H)2* +_Caso4

fl

(5) M3012 + ca(0H)2 Ms(OHi2§ + 03012

The first equation shows the reaction of calcium.(as cal-
cium bicarbonate) with lime resulting in precipitation of the
normal carbonate (CaCOs). The-removal is not complete be-
cause CaCOs is soluble to the extent of about 15 p.p.m. The
other equations show the removal of magnesium by precipita-
tion as magnesium hydroxide. If insufficient lime is added,
M3603, a soluble compound, is formed as shown in equation (2).
The use of sufficient lime results in.the formation and pre-
cipitation of Mg(0H)2 as shown in equations.(3), (4), and (5).

From these equations it will be seen that twice as much lime

(12)

is needed to remove magnesium bicarbonate as to remove cal-
cium.bicarbonate.

The lime also reacts with any free carbon dioxide (002)
present as follows:

C02 4' Ca(OH)2 :3 Ca005+ 4' H20

Lime is, therefore, required, (1) to react with all of the
bicarbonate present in the water, and (2) to react with the
magnesium in any form)aWfl/(3) 7‘0 read" “”74 71467,}? (02'

To calculate the amount of lime required for the soften-
ing reaction it is necessary to make certain chemical tests.
The tests needed are (1) free carbon dioxide (002); (2) al-
kalinity; (3) non-carbonate (incrustant) hardness; (4) cal-
cium; and (5) magnesium. These tests are made as follows:

(1) Free Carbon Dioxide

Fill a 100 m1. Nessler tube to the mark with the fresh-
ly taken sample. '

Add 10 drops of phenolphthalein indication.

Titrate rapidly with N/44 sodium hydroxide.(NaOH) from
a burette, stirring gently, until a slight permanent
pink color appears. Record the number of m1. of NaOH

used.

M1. of N/44 NaOH x 10 = p.p.m. 002

This test was not made in the course of this work because
of the fact that the sample of water stood for some time in
the laboratory. This test for 002 should be made at the time

the sample is taken because the gas escapes when the water

(13)

stands for any length of time unless the sample bottle be
kept tightly stoppered and at a temperature lower than at
which the sample was cellected.
(2) Alkalinity
Pipette 100 ml. of the sample into one Erlenmeyer
flask and the same quantity of distilled water into
another.
Add four drOps of phenolphthalein indicator to each.
If the sample becomes pink, add 0.02 N. sulphuric acid
from a burette until the pink color Just disappears
and record the number of ml. of acid used.
Add two drops of methyl orange indicator to each flask.
If the sample becomes yellow, add 0.02N sulphuric acid
until the first difference in color is noted when comp
pared with the distilled water. The end point is
orange. Record the number of ml. of acid used.
Calculations for alkalinity titrations are made as fol-
lows, giving all results in p.p.m. as CaCOs:
Let P g ml. of 0.02 N sulphuric acid used for the ti-
tration with phenolphthalein and T a the ml. of the
acid used for the tetal titration (phenolphthalein plus
methyl orange).
There are five possible conditions as follows:
1. P g T
Hydroxide (p.p.m.) = P X 10
2. P greater than i T
Hydroxide (p.p.m.) . (2P - T) X 10

Normal carbonate (p.p.m.) - 2(T - P) X 10

(14)

3eP.%T
Normal carbonate (p.p.m.) a T X 10
4. P less than % T

Normal carbonate (p.p.m.) - 2P X 10
Bicarbonate (p.p.m.) - (T 2P) X 10
5.P=°

Bicarbonate (p.p.m.) : T X 10

(3) Non-Carbonate Hardness
The carbonate and non-carbonate hardness may be calcu-
lated from the results of the alkalinity and total
hardness determinations.
Calculations-
Carbonate hardness:
Let p.p.m. normal carbonate alkalinity - p.p.m. bicar-
bonate alkalinity 3 S.
Case I. Where S is equal to or less than the total
hardness, then S a carbonate hardness.
Case II. Where S is greater than the total hardness
then the total hardness : the carbonate hardness.
Non-carbonate hardness:
P.p.m. total hardness - p.p.m. carbonate hardness .
p.p.m. non-carbonate hardness.

(4) Calcium
Evaporate to dryness, in an evaporating dish on a water
bath, 250 ml. of the well mixed sample to which a few
drOps of concentrated.bwdrochloric acid have been added.

Moisten the residue with a few drops of concentrated

(15)

hydrochloric acid.

Add 30 m1. of distilled water and heat to boiling.
Filter, rinse the dish and wash the residue on the paper
with distilled water, adding the filtered rinsings and
washings to the filtrate, but keeping the volume small.
Add a few drops of bromine water to the filtrate and
boil gently for 5 minutes.

Add 10 ml. of a 10 % solution of ammonium chloride, make
alkaline to litmus with concentrated ammonium hydroxide
and boil gently for a few minutes.

Filter, rinse the beaker and wash the precipitate on the
paper with small portions of distilled water, adding the
filtered rinsings and washings to the filtrate.

Warm the filtrate and add slowly, with constant stirring,
10 ml. of saturated ammonium oxalate.

lllow to stand for 50 minutes in a warm place.

Filter through a quantitative filter paper, rinsing the
beaker and paper with small portions of hot distilled
water. The filtrate from this filtration is saved for
the magnesium determination.

Place the original beaker from Which the filtration was
made under the funnel and, piercing a hole in the filter
paper in the funnel, wash the precipitate into the beak-
er, using about 30 ml. of 2 % (2 ml. of concentrated
acid per 100 ml.) sulphuric acid.

Heat to boiling and add from a burette a standard solu-

tion of potassium.permanganate until the first permanent

(15)

pink color is obtained.
Transfer the filter paper to the beaker and continue
the titration with potassium permanganate to the first
permanent pink color. Record the total ml. of perman-
ganate used.
M1. of KMnO4 X 10 - p.p.m. calcium (Ca).
(5) Magnesium
Use the filtrate Which was saved in the calcium deter-
mination and continue as follows-
Add an amount of concentrated ammonium hydroxide to
the filtrate equal to about 1/9 its volume and then add
10 ml. of a lO % solution of disodium.phosphate. Stir
the solution vigorously for five minutes.
Allow to stand at least four hours, and filter the pre-
cipitate and several rinsings of the beaker onto a quan-
titative filter paper, being sure that all of the pre-
cipitate is transferred to the paper.
Wash with small portions of dilute ammonia water.
Ignite, cool and weith a weigh a clean crucible.
Transfer the folded filter paper containing the magnesium
precipitate to the crucible.
Ignite until almost white, cool in a dessicator and weigh.

Gain in wt. (grams) X 875.6 g p.p.m. Mg.

(17)

After making the required analytical tests as described
above the next step in the calculation of the lime necessary
is to convert the values of the bicarbonate alkalinity and
magnesium into a common unit of measure to furthur facilitate
computations. It was decided to use the equivalent values
as p.p.m. of CaCO3. The alkalinity results were already given
as converted to p.p.m. of CaCO3 in the statement of results.

Also the molecular weight of CaCO is 100.08 amd this makes a

3
very convenient number to use in computations, as will be
seen presently.

To convert p.p.m. of Mg (as Mg) to p.p.m. of Mg (as CaCOB)
it is necessary to multiply the p.p.m. of mg (as Mg) by the
fraction whose numerator is the molecular wieght of the equi-
valent compound sought (Ca COB) and whose denominator is the
molecular weight of the known material (Mg). The molecular

weight of mg is 24.32. Hence to convert Mg (as Mg) to Mg

(as CaCO3) it is necessary to multiply the value obtained in
100.08,

24.32
After converting the Mg values to equivalent CaCOB, the

the analysis for Mg by

amounts of CaCO3 for the bicarbonate alkalinity and the Mg
are added together. This gives the total value of alkalinity
and Mg reacting with the Ca(0H) during the softening reaction.

The quantity of Ca(0H)2 required is computed in a similiar
manner as above described for converting the Mg to¢§quivalent
Ca003. The molecular weight of Ca(0H)2 is 74.10 and the
conversion fraction here becomes Iggfég. The result of this
computation will be in p.p.m. of pupa calcium hydroxide.

Since the Ca(0H)2 obtained commercially is not pure the

amount Of 08(0H)2 required by the above calculations should

(18)

100
per cent Ca(0H)? in the calcium hydrocide used
If the above calculations are carried out using p.p.m. of

 

be multiplied by

the hardness and the rest of the work is carried out in the
metric system, the resultant answer will be the amount dn grams)
of commercial calcium hydrdgide required to bring about a
theoretical 100% lime softening in a one liter sample of water.

Having calculated the amount of Ca(0H)2 required to bring
about 100% lime softening, the amounts needed for the following
percentages of the 100% treatment were computed,- 15%, 30%,
45%, 60%, 75%. 90‘, 100%, 1151, 130% and 1501.. This range
was selected as giving a sufficiently broad range of treatment
to observe the effects of the treatment.

Two samples of water with highly different qualities were
selected for this investigation.

The first water selected for this investigation was the
College tap water. This water was drawn from the college
supply pipes after allowing a sufficient quantity to run to
waste in order that the iample would be a representative one.
The College water is a ground water of average hardness,
having a total hardness of about 300 p.p.m., together with a
low chloride contentflnr/a /ola non-carboxa?€ darn/yawn

The second water was a sample of the river water from
Midland, Michigan. This water has a high total hardness,
containing nearly 800 p.p.m. of hardness, as well as a very
high cggagi/dzécgrffie-ggfbdflf high/merhloride content in the water
at Midland can be largely attributed to the presence of salt

mines and oil wells in that vicinity.

(19)

Tests were run throughout the investigation to determine
if any reduction in chloride content was obtained during
lime softening. This test was made by titrating a 50 m1.
sample of the water with a standard silver nitrate solution

and potassium chromate as an indicator.

(20)

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(23)

DISCUSSION OF GRAPES

Figure l. Alkalinities of College Tap Water During Lime
Treatment.

This graph clearly shows the removal of bicarbonate
alkalinity in proportion to the percentage of the calculated
theoretical complete lime softening. At the theoretical 100%
treatment the bicarbonate alkalinity was practically removed
and as the lime treatment was increased to excess treatment
the hydroxide alkalinity began to increase in proportion to
the increase in lime. This was due to the direct addition of

the hydroxide in the form of the lime (calcium hydroxide).

 

.(‘b'ON

.ACL‘CHL IN'O.

(25)

Figure 2. Effect on Calcium and magnesium Content of
College Tap Water during Lime Treatment.
This graph shows that the calcium in the water is removed
at a fairly constant rate until it reached its minimum at the
point where the treatment was 100% of the theoretical. After
the point mentioned above, the lime treatment was excessive
and the calcium content rose rapidly as the lime treatment was
increased, {/up 7’0 7’1“: Iin’ror/uc/z‘ou of Jo/z/A/e m/cu'zm Ay%o/I'q@
The magnesium content remained practically uniform until
the lime treatment began to be in excess of that required for
the W. This shows that the magnesium requires ”7
more lime for its removal than does calcium. It also shows
that magnesium is not removed by lime treatment until pract-
[1:1 749 form 07‘ b/‘cr/r/aoka
ically all the calcium which~eill~beAremevedebyalimeetreatment
is precipitated out of solution. 'Upon the addition of suf-
ficient lime practically all the magnesium is removed. This

is effected at about 145% of the theoretical treatment.

(25)

 

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(27)

Figure 3. Total Hardness, Carbonate Hardness and Non-
Carbonate Hardness of College Tap Water During
Lime Treatment.

These graphs show the presence of free 002 in the original
sample as drawn from the tap and its removal by the lime at the
beginning of the lime treatment.

The removal of total hardness and the carbonate hardness go
parallel to each other until the total hardness begins to
increase due to the increase in calcium as shown in Figure 2.
The carbonate hardness continues to decrease due to the
beginning of the magnesium removal.

The non- carbonate curve shows a decided increase at the 90%
treatment. This is ppparently due to some experimental error 5?
as the curve drops sharply to about where it was before.

Then as the hydroxide alkalinity (see Pigurefil) increased due
to the excess of calcium hydroxide treatment the non-carbonate
hardness showed a corresponding increase. This shows the

non-carbonate hardness to be due chiefly to hydroxide alkalinity

caused by the calcium hydroxide.

(28)

 

IOIYON

'NC

IACLA'NI flh ..

(29)

Figure 4. Alkalinities of Midland River Water During
Lime Treatment.

This graph is very similiar to Figure 1 showing a rapid
reduction in bicarbonate alkalinity until completely removed
and immediately following the complete removal of the bicar-
bonate alkalinity, the hydroxide alkalinity began to increase
as the lime dosage increased.

The chief difference between this graph and that shown in
Figure l is that in the latter the bicarbonate alkalinity
was removed at the 100% theoretical lime treatment point for
the College water while in Figure 4, the zero for bicarbonate
alkalinity occurs at 75% of the theoretical treatment necess-
ary for the Midland water. This may be explained by the fact
that in the College water the principal hardness is caused
by the bicarbonate alkalinity while in the Midland water the
bicarbonate alkalinity is only about half that of the College
water, but the calcium is three times as great and the mag-
nesium is twice as great in the Midland water as in the

College water.

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(31)

Figure 5. Effect on Calcium and Magnesium Content of
Midland River Water During Lime Treatment.

These graphs show even better than Figure 2, the order in
which the calcium and magnesium are removed in the lime
softening process. Here there is no lag in.the return of the
calcium upon the continued lime treatment. In this case the
magnesium is being removed very slowly while the calcium is i
being precipitated. As soon as the calcium reaches a maximum
of removal and begins to increase due to the excess treatment
the rate of removal of the magnesium increases and continues
until the magnesium is almost entirely removed at 145%
treatment point.

(32)

 

(33)

Figure 6. Total Hardness, Non-carbonate Hardness and
Carbonate Hardness of Midland River Water During
Lime Treatment.

These graphs are comparable with those in Figure 3 and the
same explanation holds true here as diJ in that case. The
non-carbonate hardness shows some fluctuation here, but not a
sharp increase as it did.with the College water in Figure 3.
This gives credence to the explanation of its nise there as
being due to experimental error rather than some behavior in
the softening reéiion. Altogether these curves show the
paralleling of the hardness relationships during the lime

treatment.

(34')

 

IACLACNI Oh I. (NC Iin'ou

 

Figure 7. Comparison of Total Hardness - Midland River
Water and College Tap Water During Lime Treatment.

These graphs have no special significance except as they
show that in both cases of the treatment of such radically
different waters as were used in this investigation the amount
of change in the total hardness was about the same,.being about
200 p.p.m. In both cases the continued treatment with lime
caused an increasefin‘total hardness. This increase has been
noted in the previous discussion and will not need to be
discussed again here. It should be noted, however, that the
lowest amount of total hardness is reached earlier in.the
Midland water treatment than in the treatment of the College
water. This has also been discussed before (see discussion

of Figure 4), but is again shown.here.

CONCLUSION

From the foregoing discussion and from my results of the
treatment of these two different waters with calcium hydroxide
I feel that it is possible to draw a few general conclusions.

It must be kept in mind that there is danger in forming
too specific a set of conclusions from such a limited number
of cases. If time had permitted, I would like to have been
able to carry on the investigation of these softening reattions
so as to have included a greater number of kinds of water.
This would have enabled me to draw a more complete and a truer
picture of what actually can be expected of the lime softening
process. Time would not permit carrying out the great number
of anéyses necessary so it was necessary to select a ground
water of average (for Michigan) hardness and a river water of
a very high degree of hardness for my investigation.

In interpreting the results of these analyses it must also
be borne in mind that no matter how carefully one carries out
an analysis a certain amount of experimental and personal
error is bound to creep into the work. The magnitude of such
errors is indeterminate, but based on the laws of probability
they may be assumed to be small. Accordingly, small
deviations from an average, or an anticipated, value must not
be given too much importance.

With the abate limitations in mind I have drawn the follow-
ing conclusions.

1. Lime softening alo does not remove any appreciable

amount of iron or aluminu/

(37)

2. Lime softening alone does not remove any of the chlorides

fill/'0 {/5/\

”02’ 7/,

_’

present in the water. The presence of chlorides apparently a
does not have any affect upon the softening of the water wiéhx,
lime. This is what would be expected from a study of the
chemical equations of the reactions of lime and chlorides.
The equations point out the need for the addition of sodium
carbonate to effect the removal of chlorides.

3. Calcium is removed from the water by lime softening up
to the point where the bicarbonate alkalinity is removed.
After this point the calcium content of the water increases
with the increased addition of the calcium hydroxide being
used for the softening.

4. Magnesium is apparently unaffected until a point is re
reached Just before the calcium is most completely removed.
It then begins to be removed, quite rapidly at first and then
more slowly as it approaches the point of complete removal.

5. Carbonate hardness is removed quite rapidly until the
calcium content reaches a minimum. After this point the
carbonate hardness removal proceeds more slowly. This indicates
that the principal carbonate hardness of the waters invest-
igated was that due to calcium bicarbonates.

6. The non-carbonate hardness of the waters investigated
did not seem to bear much of a general relation to each other
until the treatment with the calcium hydroxide had reached the
point where the calcium removal was the greatest. After this
point the nonpcarbonate hardness was the kind of hardness in
which the addition of hydroxide alkalinity, by continued
treatment with the lime, was indicated. This is shown by the

tests for alkalinity.

)

(38)

Perhaps the most significant fact brought out by this
investigation as a whole is that it clearly shows that it is
impossible to adopt any "out and dried" treatment to be
followed in planning or designing a water softening plant.
Each locality has its own pecularities which it imparts to its
water supply and each water supply must be individually

considered When its treatment is being contemplated.

BIBLIOGRAPHY

BABBITT & DOLAND Water Supply Engineering

CHRISTIE, W. P. Water— Its Purification and Use

ELDRIDGE & THEROUX A Laboratory Manual for the Chemical
Analysis of Water and Sewage

HORWOOD, M. P. Sanitation of Water Supply

RIDEAB, SAMUEL Water and Its Purificat ion

STEIN, M. F. Water Purification Plants and Their’Operation
HOOVER, C. P. "Water Softening“ Ehg. Cont. Mar. 14, 1923
STEIN, M. F. “Water Softening for Muncipalities“ Jour. Am.

Water Works Assn. Vol: 6 - 1919

STEPHENSON, F. H. & Weston, R. s. "The Springwells Filtration
.Plant, Detroit, Michigan Proceedings of Am.
Soc. Civil Eng. Jan. 1935

QSoftening of Public Water Supplies" Jour. Am. Water Works Assn.

' Vol: 12 - 1924; Vol: 14 - 1925

“Water Softening & Purification at GOIUmbus, Ohio“ Pub. Works
V61? 51 - 1921 '

“Why Water Should be Treated" Domestic Engineer - Nov. 1934

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