Alkalinity and Hardness teri maaka bhosda

pH, Alkalinity and Hardness
S. Y. B. Tech (Civil)
Environmental Engineering CE234
Prof Mrs N G Gogate

pH
• Measure of Intensity of acid / alkaline condition of a solution
• Way of expressing hydrogen ion concentration OR hydrogen-ion activity.
• Importance:
• Water supply: pH needs to be considered in Chemical Coagulation, Disinfection, Water Softening
and Corrosion control
• Wastewater treatment: pH needs to be considered in chemical treatment processes used to
coagulate wastewaters, dewater sludges or oxidize certain substances (cyanide). Also, in biological
treatment processes, pH must be controlled within a range favourable to the specific organisms
involved.
• 𝐻2𝑂 ↔ 𝐻+
+ 𝑂𝐻−
and
𝐻+ ∗[𝑂𝐻−]
[𝐻2𝑂]
= 𝐾
• But, concentration of water is very large (as it gets ionized by a very small amount) as
compared to the concentration of ions, hence
• 𝐻+ ∗ 𝑂𝐻− = 𝐾𝑤 and for pure water at 25℃; 𝐻+ ∗ 𝑂𝐻− = 10−14
• When an acid is added to water, it ionizes thus increasing 𝐻+
concentration.
Consequently, the 𝑂𝐻− concentration must decrease in conformity with the ionization
constant.
• Important to remember that 𝐻+
or 𝑂𝐻−
concentration can never be reduced to zero,
no matter how acidic or basic the solution may be.

pH concept and Measurement of pH
• Expression of 𝐻+
ion concentrations in terms of molar concentrations
being rather cumbersome, pH is expressed as the negative logarithm of the
𝐻+
ion concentration.
• 𝑝𝐻 = − log[𝐻+
]
• pH scale usually represented as ranging from 0 to 14.
• Measurement of pH
• pH meter: Based on potentiometric measurement of hydrogen ion
concentration in the solution. It consists of a glass electrode (which shows
potential proportional to the [𝐻+]) and a reference electrode (having fixed
& known potential). Needs calibration with a solution of known pH.
• BIS standards
Authority HDL (Highest Desirable Level) MPL (Maximum Permissible Level)
BIS 6.5-8.5 No relaxation

Alkalinity and forms of alkalinity
• Alkalinity is:
 the ability of a solution to neutralize acids.
 the capacity of water to resist changes in pH that would make the water more acidic. (It should not be
confused with basicity which is an absolute measurement of only [𝑂𝐻−
] on the pH scale)
 the buffering capacity of water.
 a measure of acid neutralizing capacity of water.
 a measurement of dissolved alkaline substances in water.
 ability to resist changes in pH upon the addition of acids (𝐻+
ions).
• pH measures the concentration of [𝑂𝐻−] ions indirectly while alkalinity measures the capacity to
neutralize acids due to the presence of ions such as [𝐻𝐶𝑂3
−
], [𝐶𝑂3
−−
], and [𝑂𝐻−].
• Alkalinity is chiefly caused by bicarbonates, carbonates and hydroxides of alkali earth metals Ca,
Mg, K, and Na.
• Alkalinity of water is due to the presence of mainly 3 anions, [𝐻𝐶𝑂3
−
], [𝐶𝑂3
−2
], and [𝑂𝐻−].
• Hence, there are 3 forms / types of alkalinity
• Bicarbonate [𝐻𝐶𝑂3
−
]
• Carbonate [𝐶𝑂3
−2
], and
• Hydroxyl [𝑂𝐻−
].

Natural Sources of Alkalinity
• The alkalinity of natural water is determined by the soil and bedrock through which it passes. The main
sources for natural alkalinity are rocks which contain carbonate, bicarbonate, and hydroxide compounds.
• Variations in the alkalinity of inland waters can be attributed to the amount of weathering of bedrock
material and soils derived from the bedrock.
• Alkalinity of natural waters is due primarily to the presence of weak acid salts although strong bases may
also contribute (i.e. 𝑂𝐻−
) in extreme environments.
• Bicarbonates of Ca, Mg and Na are most predominant in natural underground water bodies.
(CO2 + H2O) + CaCO3 or Ca (HCO3)2 or
MgCO3 or Mg (HCO3)2 or
Na2CO3 2NaHCO3
• In natural surface water bodies, such as rivers and lakes, CO2, a product of microbial oxidation of organic
matter and vegetative respiration is removed by algal photosynthesis and surface aeration.
• This results in conversion of a part of bicarbonate alkalinity into carbonate alkalinity and even into hydroxide
alkalinity, if there is an algal bloom. pH also increases during this process and may go upto 10-11.
2𝐻𝐶𝑂3
−
→ 𝐶𝑂3
−2
+ 𝐻2𝑂 + 𝐶𝑂2 ↑ → 2𝑂𝐻− + 𝐶𝑂2 ↑
• Surface waters contain carbonate and hydroxyl alkalinity predominantly and bicarbonate to a smaller
extent.
• While groundwater predominantly contains bicarbonate alkalinity.

How minerals like Bicarbonates (𝐻𝐶𝑂3
−
), Carbonates (𝐶𝑂3
−2
), Calcium (𝐶𝑎+2
), Magnesium (𝑀𝑔+2
)
enter in water
𝐻𝐶𝑂3
−

pH and Forms of Alkalinity
• pH of water affects the relative concentrations of various forms of alkalinity.
• At low pH values, bicarbonate alkalinity is predominant, while at higher pH
values, alkalinity forms change to carbonate and hydroxyl alkalinity.
• Total alkalinity is sum of all the three forms of alkalinity.
pH Forms of Alkalinity
>8.3 >11.30 OH- only
10.57-
11.30
(OH- + CO3
- -) predominantly
or CO3
- - only
8.30-10.56 (CO3
=+HCO3
-) predominantly
or CO3
-2 only
≤
8.3
4.50-8.30 HCO3
- predominantly
<4.5 Nil

Significance of Alkalinity
• Alkalinity is important during coagulation. It buffers the water in the pH range
most suitable for effective coagulation.
• Alkalinity is important in water softening. It is used for calculating the quantities
of chemicals required such as lime and soda ash.
• Alkalinity is important in corrosion control. Slightly alkaline waters are less
corrosive then acid waters in the distribution system.
• Alkalinity determination is also important in maintaining high efficiency of
anaerobic digesters and also in determining suitability of wastes and wastewaters
for biological treatment.
• When 𝑂𝐻− alkalinity is in excess of 0.8 mg/L as CaCO3 water becomes highly
unpalatable because of caustic taste. Bleeding of the tongue and gullet may occur
if consumed. Water will be highly corrosive.
• High concentration of OH- ions replace F- ions in the insoluble coating of the
teeth. The dental enamel dissolves in acids and teeth develop cavities.

Unit of Alkalinity
• Most of the water quality parameters are expressed in mg/L (i.e. mg concentration of the
substance in 1 liter of water).
• As 3 different forms of alkalinity exist, concentrations need to be expressed as equivalent
concentrations. This facilitates conversion between total and individual forms of alkalinity.
• Alkalinity is measured in mg/L as 𝐶𝑎𝐶𝑂3.
• Concentration of various substances can be expressed in terms of 𝐶𝑎𝐶𝑂3 using following
equation.
Conc. of ‘A’ in mg/L as 𝐶𝑎𝐶𝑂3 =
𝐴𝑐𝑡𝑢𝑎𝑙 𝑐𝑜𝑛𝑐.𝑜𝑓 𝐴
𝑒𝑞𝑢𝑖𝑣𝑎𝑛𝑡 𝑐𝑜𝑛𝑐.𝑜𝑓 𝐴
∗ 𝑒𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝑜𝑓 𝐶𝑎𝐶𝑂3
• Equivalent concentration of 𝐶𝑎𝐶𝑂3 = 50 eq/L

Determination of Alkalinity
• Measured by titrating the sample with 0.02N 𝐻2𝑆𝑂4.
• 1 ml of 0.02 N H2SO4 = 1 mg of alkalinity as CaCO3
• For determining forms of alkalinity, 2 indicators (Phenolpthalein (PP) and Methyl Orange (MO) are
used.
• If pH≥ 8.3, PP indicator is used followed by MO, else directly MO indicator is added.
• Graph shows change in pH with addition of acid. 2 inflection points can be seen, first at pH=8.3 and
second at pH=4.5. Below pH=4.5, there is no alkalinity.
• mL of 𝐻2𝑆𝑂4 till PP end point ‘x’; and mL of 𝐻2𝑆𝑂4from PP to MO end point ‘y’
𝑇𝐴 =
𝑥+𝑦 ∗0.02∗50∗1000
𝑠𝑎𝑚𝑝𝑙𝑒 𝑣𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝑚𝐿
,
𝑚𝑔
𝐿
𝑎𝑠 𝐶𝑎𝐶𝑂3
When the pH of the sample is 8.3, the amount of titrant used (‘x’ ml) up to pp end point is a measure
of
 All the OH- alkalinity, or
 All the OH- alkalinity + ½ of CO3
= alkalinity or
 ½ of CO3
= alkalinity only.
The amount of titrant used between PP end point and MO end point (‘y’ mL) is a measure of
 1/2 of the Co3
= alkalinity only or
 1/2 of CO3
= alkalinity + all the HCO3
- alkalinity or
 only HCO3
- alkalinity (if the pH of the sample at the start of titration is 8.3 or less).

Determination of forms of alkalinity
• The reactions at the end points are-
• 𝑂𝐻− + 𝐻+ → 𝐻2𝑂; and
• 𝐶𝑂3
−2
+ 𝐻+ → 𝐻𝐶𝑂3
−
at PP end point (pH=8.3)
• 𝐻𝐶𝑂3
−
+ 𝐻+
→ 𝐻2𝐶𝑂3 at MO end point (pH = 4.5)
• Case—1:-
• (only x present, Y=0)
• x=ml of titrant used up to PP end point
• y= ml of titrant used between PP end point and MO end point.
• Alkalinity present only OH-
• 𝑇𝑜𝑡𝑎𝑙 𝑎𝑙𝑘. = 𝑂𝐻− 𝑎𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦,
𝑚𝑔
𝐿
𝑎𝑠 𝐶𝑎𝐶𝑂3 =
𝑥 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)
𝑆𝑎𝑚𝑝𝑙𝑒 𝑡𝑎𝑘𝑒𝑛 (𝑚𝐿)

• Case 2: - (x > y)
• Alkalinity present (OH- + CO3
=)
• 𝑇𝑜𝑡𝑎𝑙 𝑎𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦,
𝑚𝑔
𝐿
𝑥+𝑦 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)
• 𝐶𝑂3
−2
𝑎𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦,
𝑚𝑔
𝐿
2𝑦 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)
• 𝑂𝐻− 𝑎𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦,
𝑚𝑔
𝐿
𝑎𝑠 𝐶𝑎𝐶𝑂3 = 𝑇𝑜𝑡𝑎𝑙 𝑎𝑙𝑘. − 𝐶𝑂3
−2
alk.
• Case 3: - (x=y)
• Alkalinity present – CO3
= only
• 𝐶𝑂3
−2
𝑚𝑔
𝐿
2𝑥 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)

• Case 4: - (x < y)
• Alkalinity present (CO3
= + HCO3
-)
• 𝐶𝑂3
−2
𝑚𝑔
𝐿
2𝑥 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)
• 𝑇𝑜𝑡𝑎𝑙 𝑎𝑙𝑘𝑎𝑙𝑖𝑛𝑖𝑡𝑦,
𝑚𝑔
𝐿
𝑥+𝑦 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)
• 𝐻𝐶𝑂3
−
𝑚𝑔
𝐿
𝑎𝑠 𝐶𝑎𝐶𝑂3 = 𝑇𝑜𝑡𝑎𝑙 𝑎𝑙𝑘. − 𝐶𝑂3
−2
alk.
• Case 5:- x = 0, only y present
• Alkalinity present, only HCO3
-
• 𝑇𝑜𝑡𝑎𝑙 𝐴𝑙𝑘. = 𝐻𝐶𝑂3
−
𝑚𝑔
𝐿
𝑦 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)

BIS Standards (IS 10500:2012)
Acceptable Limit (mg/L as 𝑪𝒂𝑪𝑶𝟑) Permissible Limit (mg/L as 𝑪𝒂𝑪𝑶𝟑)
200 600
• Low Alkalinity (i.e. high acidity) causes deterioration of plumbing and increases the
chance of dissolution of many heavy metals in water that are present in pipes, solder or
plumbing fixtures.
• Beyond the acceptable level of alkalinity, taste becomes unpleasant.
• Beyond the permissible limit, high concentration of OH- ions replace F- ions in the
insoluble coating of the teeth. The dental enamel dissolves in acids and teeth develop
cavities.
• No standards for minimum desirable alkalinity have been specified for drinking water.
However, bicarbonate alkalinity of about 75 mg/l (not less than 50 mg/l) as CaCo3 may
be considered desirable which will contribute to the wholesome mineral content of
drinking water.
• When OH alkalinity is in excess of 0.8 mg/l as CaCO3 water becomes highly unpalatable
because of caustic taste bleeding of the tongue and gullet may occur if consumed. Water
will be highly corrosive.

Hardness
As water moves through soil and rock, 𝐶𝑂2 (a product microbial action on organics)
dissolves in it making it slightly acidic. This slightly acidic water dissolves very small
amounts of minerals and holds them in solution.
Calcium and magnesium dissolved in water are the two most common minerals that
make water "hard." The degree of hardness becomes greater as the calcium and
magnesium content increases.
Hardness prevents soap from lathering. Earlier, hard waters were identified based on
their property of not forming lather with soap.
Hardness is a measure of 𝑪𝒂+𝟐 𝒂𝒏𝒅 𝑴𝒈+𝟐 ion concentration in water.
Other multivalent ions like 𝑆𝑟+2
, 𝐹𝑒+3
, 𝑀𝑛+2
also impart hardness (but rarely present).
• Significance of Hardness
Hard water causes scales to build up in boilers, as well as household appliances and fixtures.
While hard water results in high soap usage, soft waters need excessive amounts of water for
washing the lather.
𝐶𝑎+2
𝑎𝑛𝑑 𝑀𝑔+2
ions help in enzymatic processes in heart muscles. It is statistically observed that
cardiovascular diseases are prevalent in regions where people drink soft water as compared to
where people drink harder water.
Moderately hard water is suitable for drinking.

Types of Hardness
• Carbonate / Temporary Hardness (CH): Hardness caused due to bicarbonates and
carbonates (𝐻𝐶𝑂3
−
𝑎𝑛𝑑 𝐶𝑂3
−2
) of Calcium and Magnesium (𝐶𝑎+2 𝑎𝑛𝑑 𝑀𝑔+2).
This hardness can be removed using simple methods such as boiling, thus called
as temporary.
• Non-carbonate / Permanent Hardness (NCH): Hardness caused due to Chlorides
(𝐶𝑙−
), Sulphates (𝑆𝑂4
−2
), Nitrates (𝑁𝑂3
−
) of Ca, Mg, Sr, Fe, Mn. Softening
methods like Lime-soda, Ion exchange, etc. are required for removing this
hardness, hence called as Permanent.
• Unit of measurement: mg/L as 𝐶𝑎𝐶𝑂3 (same as alkalinity). This helps in
performing numerical operations on hardness and alkalinity.
• How to calculate CH and NCH using (𝐻𝐶𝑂3
−
+ 𝐶𝑂3
−2
) alkalinity
• Let total hardness of a water sample be denoted by TH; (𝐻𝐶𝑂3
−
+ 𝐶𝑂3
−2
)
alkalinity by A.
• 𝐼𝑓 𝑇𝐻 > 𝐴; 𝑡ℎ𝑒𝑛 𝐶𝐻 = 𝐴 𝑎𝑛𝑑 𝑁𝐶𝐻 = 𝑇𝐻 − 𝐴
• 𝐼𝑓 𝑇𝐻 < 𝐴; 𝑡ℎ𝑒𝑛 𝐶𝐻 = 𝑇𝐻 𝑎𝑛𝑑 𝑁𝐶𝐻 = 0.

BIS Standards (IS 10500:2012)
Acceptable Limit (mg/L as 𝑪𝒂𝑪𝑶𝟑) Permissible Limit (mg/L as 𝑪𝒂𝑪𝑶𝟑)
200 600
• Whereas hard water is inconvenient for cooking and washing soft water is unfit
for drinking. As a compromise, moderately hard water is considered desirable for
domestic supply.
• Indian standards set higher permissible levels as the majority of Indian Population
being in rural areas chiefly depends upon shallow and deep wells for their water
requirement and lower permissible values would warrant softening of drinking
water which would lay undue stress on treatment facilities. This is based on the
fact that there is no evidence, that drinking hard water is harmful to health.
• Total hardness in excess of the maximum permissible level causes excessive scale
formation and when consumed it can have a laxative effect particularly in the
case of un-acclimatized individuals.

Determination of hardness
• It is determined by titrating the sample with 0.02N EDTA, using Eriochrome Black
T as indicator.
• Principal of test:
Eriochrome black T, an organic indicator dye, forms a sky blue colored solution when this is
added to a sample of hard water at pH 10. This dye forms a weak wine red complex with
hardness producing cations, principally Ca2+, Mg2+ .
To this complex, if EDTA is added, EDTA extracts Ca2 ions to form a stable complex. When all
the Ca2+ and Mg2+ ions are extracted, the indictor is totally released in its natural sky blue
colour at the end point. The amount of EDTA used (‘x’ mL) is a measure of the concentration
of Ca2+, Mg2+ and other divalent ions causing hardness.
• A blank titration using distilled water is done following the same procedure. The
amount of titrant required for blank titration is ‘y’ mL.
• 𝑇𝑜𝑡𝑎𝑙 𝐻𝑎𝑟𝑑𝑛𝑒𝑠𝑠,
𝑚𝑔
𝐿
𝑥−𝑦 𝑚𝐿 ∗0.02
𝑚𝑔
𝑚𝐿
∗50 ∗1000 (
𝑚𝐿
𝐿
)

Calculating Equivalents
• When ions or radicals (compounds) react with each other to form
new compounds, the reactions may not proceed on a one-to-one
basis (Example 𝑁𝑎𝐶𝑙 ↔ 𝑁𝑎+
+ 𝐶𝑙−
).
• Many reactions proceed on an equivalence basis that can be related
to electro-neutrality.
• Equivalence of an element or radical is defined as the number of
hydrogen atoms that element/radical can hold in combination or can
replace in reaction (In most cases, the equivalence of an ion is same
as the absolute value of its valence.
• An equivalent of an element is its gram molecular mass divided by its
equivalence. (𝑀𝑖𝑙𝑙𝑖𝑒𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡 =
𝑚𝑎𝑠𝑠 𝑖𝑛 𝑚𝑔
𝑒𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑐𝑒
)

How many grams of 𝐶𝑎+2
will be required to combine with 90 grams of
𝐶𝑂3
−2
to form 𝐶𝑎𝐶𝑂3
Equivalent of 𝐶𝑂3
−2
= [12+(3*16)]/2= 30 g/equiv.
Equivalent of 𝐶𝑎+2 = 40/2= 20 g/equiv.
Number of equivalents of 𝐶𝑎+2 must equal the number of equivalents
of 𝐶𝑂3
−2
.
No. of equivalents of 𝐶𝑂3
−2
= 90 g/(30 g/equiv.) = 3
Hence, 3 equivalents of 𝐶𝑎+2
OR 3 equiv.*20 g/equiv. = 60 grams of
𝑪𝒂+𝟐
will be required to react with 90 grams of 𝐶𝑂3
−2
.

Equivalents also provide means of expressing various constituents of dissolved solids in a
common term. An equivalent of one substance is chemically equal to an equivalent of any
other substance. Thus, concentration of substance ‘A’ can be expressed as an equivalent
concentration of substance ‘B’ using the following relation.
𝑔
𝑙
𝐴
𝑔
𝑒𝑞𝑢𝑖𝑣
𝐴
∗
𝑔
𝑒𝑞𝑢𝑖𝑣
𝐵 =
𝑔
𝑙
𝐴 expressed in terms of ‘B’
What is the equivalent 𝐶𝑎𝐶𝑂3 concentration of (a) 117 mg/l of NaCl and
(b) 2 ∗ 10−3 moles of NaCl?
(a) Equivalent of NaCl = (23+35.5)/1 = 58.5 g/equiv.
Equivalent of 𝐶𝑎𝐶𝑂3 = (40+12+(3*16))/2 = 50 g/equiv.
NaCl concentration expressed as 𝐶𝑎𝐶𝑂3 = (117 mg/l) / (58.5 mg/mequiv.) * 50
mg/mequiv. 𝐶𝑎𝐶𝑂3 = 100 mg/l as 𝐶𝑎𝐶𝑂3
(b) NaCl concentration as mg/l of 𝐶𝑎𝐶𝑂3= [NaCl (moles/l) /NaCl (moles/equiv.)]*
𝐶𝑎𝐶𝑂3 (g/equiv.)
=
2∗10−3
1
∗ 50 = 0.1
𝑔
𝑙
= 100 𝑚𝑔/𝑙

Homework
1. If 𝐾𝑤 is 1.5 ∗ 10−15
at 10℃, what is the pH of pure water at 10℃?
2. The pH of a solution is 9.1. What is the concentration of hydroxide ions in
this solution?
3. How many grams of 𝐶𝑎𝑂 are required for 246 g 𝑀𝑔(𝐻𝐶𝑂3)2?
4. Express following concentrations of elements and compounds as mg/l of
𝐶𝑎𝐶𝑂3:
a) 95 mg/l 𝐶𝑎+2
b) 420 mg/l 𝑀𝑔𝑆𝑂4
c) 87 mg/l 𝑀𝑔+2
d) 189 mg/l 𝑁𝑎𝐻𝐶𝑂3
5. Express following molar concentrations as mg/l of 𝐶𝑎𝐶𝑂3.
a) 1 ∗ 10−2 𝑚𝑜𝑙𝑒𝑠
𝑙
𝑜𝑓 𝐴𝑙+3
b) 1.8 ∗ 10−3 𝑚𝑜𝑙𝑒𝑠
𝑙
𝑜𝑓 𝐶𝑎𝑆𝑂4

Ion Balance (Total Dissolved Solids)
• The ions usually accounting for the vast majority of TDS in natural
waters are
𝑁𝑎+, 𝐶𝑎+2, 𝑀𝑔+2, 𝐻𝐶𝑂3
−
, 𝑆𝑂4
−2
, 𝐶𝑙−, 𝐹𝑒+2, 𝐾+, 𝐶𝑂3
−2
, 𝑁𝑂3
−
, 𝐹−
• Out of these, some ions which are often sufficient to characterize the
dissolved solids content of water, are measured individually and
summed on an equivalent basis to represent the approximate
concentration of TDS. As a check, the sum of the anions should equal
the sum of the cations because electro-neutrality must be preserved.

Problem: Tests for common ions are run on a sample of water and the
results are given below. Draw a bar diagram and calculate percent
error. 𝐶𝑎+2 = 55
𝑚𝑔
𝑙
; 𝐻𝐶𝑂3
−
= 250
𝑚𝑔
𝑙
; 𝐶𝑙− = 89
𝑚𝑔
𝑙
; 𝑀𝑔+2 =
18
𝑚𝑔
𝑙
; 𝑆𝑂4
−2
= 60
𝑚𝑔
𝑙
; 𝑁𝑎+ = 98 𝑚𝑔/𝑙
(Answer: % error = 8.27%)
Cation
name
Concentra
tion
(mg/l)
Equivalent
weight
(mg/mequiv.
)
Equivalent
concentratio
n (meq/l)
Anion
name
Concen
tration
(mg/l)
Equivalent
weight
(mg/mequiv.)
Equivalent
concentration
(meq/l)
𝐶𝑎+2 55 20 (55/20) =
2.75
𝐻𝐶𝑂3
−
250 (61/1)=61 (250/61)=4.1
𝑀𝑔+2 18 (24/2)=12 (18/12)=1.5 𝑆𝑂4
−2 60 (96/2)=48 (60/48)=1.25
𝑁𝑎+ 98 (23/1)=23 (98/23) =
4.26
𝐶𝑙− 89 (35.5/1)=35.5 (89/35.5)=2.51
Total 8.51 meq/l Total 7.86 meq/l

Example
The ion concentration obtained for a groundwater sample is as follows. Determine
alkalinity, total hardness, carbonate and non-carbonate hardness. 𝐶𝑎+2 = 180mg/L;
𝑀𝑔+2
= 48 mg/L; 𝐻𝐶𝑂3
−
= 183 mg/L; 𝐶𝑂3
−2
= 180 mg/L; 𝑆𝑂4
−2
= 40 mg/L; 𝐶𝑙−
=
15 mg/L.
Alkalinity is caused due to bicarbonates, carbonates and hydroxyl ions. For this
sample, hydroxyl ion concentration is zero, thus
Total Alkalinity=bicarbonate + carbonate alkalinity.
Bicarbonate and carbonate ion concentrations have to be converted in terms of
𝐶𝑎𝐶𝑂3 and then adding these concentrations, total alkalinity can be calculated.
Total Alkalinity =
183
61
+
180
30
∗ 50 = 450
𝑚𝑔
𝐿
Equivalent weight of 𝐶𝑂3
−2
= 30 and Equivalent weight of 𝐻𝐶𝑂3
−
=61 eq/L
Total Hardness =
180
20
+
48
24
∗ 50 = 550
𝑚𝑔
𝐿
Total Hardness> (Bicarbonate +carbonate) alkalinity, CH =450
𝑚𝑔
𝐿
NCH= TH-CH= 550-450= 100
𝑚𝑔
𝐿

Alkalinity and Hardness teri maaka bhosda

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