Sampson 1
Dillon Sampson
Chem 111
Section 104
April 3, 2012
Determining the Hardness of Water Using Multiple Techniques
Group Members:
Dillon Sampson
Kyle Sam
Sydney Sherman
Mike Shaw
TA: Nick McCool
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Introduction:
Water Hardness refers to the concentration of Mg2+ and Ca2+ along with other
polyvalent cations in a water sample. Water increases in hardness as it picks up magnesium and
calcium cations from the earth (1). Due to its polarity water has an overall negative charge. This
charge is what attracts (or picks up) these positively charged magnesium and calcium cations and
their compounds. In theory, water hardness increases as the amount of contact with soil and rock
increases because more of these cations are taken up (dissolved). The United States Geological
Survey (USGS) on water quality gives a scale on the hardness of water. Usually water hardness
is measured in terms of the levels of calcium carbonate. A water sample with a greater amount
of calcium carbonate (by concentration) would be considered harder, and softer with less
concentration. The USGS scale puts the water into categories of soft, moderately hard, hard, and
very hard based on the calcium carbonate concentration (2,3). While it has little affect on health,
the hardness of natural waters is relevant when dealing with washing processes, plumbing, and in
industrial boilers and evaporators. Calcite deposits which are known as scales form that can
cause various problems including energy inefficiency, blockages, and expenses for removal
(1,4).
There are multiple ways to measure water hardness including TDS (total dissolved
solids), EDTA (ethylenediaminetetraacetic acid) titration, and AA (atomic absorption
spectrophotometry). The presence of cations in water is commonly measured in parts per million
(12). EDTA titration uses a complexation reaction to determine quantitatively the total amount
of divalent cations in a sample of water using ethylenediaminetetraacetic acid as a chelating
agent. The EDTA chelating agent along with an NH3/NH4 buffer and a high pH indicator
(erichrome black T, or EBT) are used in the titration. The EDTA reacts with the cations to
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basically remove them from the sample and soften the water. The buffer is used to set the pH to
about 10 in order for the blue form to predominate. Once EBT is added it should be in its blue
form. More drops of EDTA required to turn the water blue meant the water sample was harder.
The harder water had a higher number of ions and therefore required more EDTA to be fully
chelated. The point in the well set where the color changes from red to blue is the point at which
full chelation has occurred. Using the amount of EDTA to fully chelate the sample ions, the
number of ions in the sample can be determined. AA is another way to determine the
concentration of ions in a sample, but is much different than EDTA titration. AA uses a light
source set to the wavelength and energy of the atoms in question, Mg2+ and Ca2+ in this case,
which is shined through the sample being burned inside the instrument. The light absorption
level is proportional to the ion concentration of the water sample. An AA instrument is
calibrated to the specific targeted ion under observation. Only one absorbance level can be
determined at a time (Mg2+ value separately obtained from the Ca2+ value) (1). While EDTA
provides values for concentration, AA gives values for light absorption which must be converted
over to molarity. AA offers a more consistent procedure, but it is not necessarily more accurate.
The biggest difference in AA and EDTA titration is the technology involved with AA. The
addition of Baking Soda (NaCO3) and cation exchange resin were the two methods for softening
the sample. In both cases cation exchange occurs where the divalent cations (Mg2+ and Ca2+)
take the place of the softer monovalent cations (Na+ and H+) in the commercial water softener
and the resin. The sample of water taking on the softer monovalent cations then becomes
“softer” (5).
Water samples in State College taken from the first floor of the HUB (6), the second floor
of the HUB (7), the third floor of the HUB (8), and the filtration station on the second floor of
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the HUB (9) were analyzed to determine hardness. State College receives water from the same
ground water supply. The State College water is relatively hard with calcium concentration
between 180-220 parts per million (10). The only variability in samples taken across State
College should be from filtration setups and aged piping. Since our samples were taken from the
same building the piping should play no factor in the hardness of each sample. It is predicted
that the samples taken from the three floors will have equal hardness values (divalent ion
concentrations) because they are from the same source and flow through the same piping system.
The water sample taken from the filtration station will be softer because the system removes
impurities.
Procedure:
In order to test differences in water from State College the four samples had to be taken.
Three of the samples were taken from taps located in the bathrooms of the designated floors
(1st,2nd, and 3rd). The fourth sample was taken from a filtered water tap on the second floor.
The PSU Chemtrek Lab Manual provided the procedure for going about testing the hardness of
the samples. Using the AA spectroscopes provided in lab, atomic absorbance values were
determined for Mg2+ and Ca2+. The concentration values in absorbance were recorded along
with the AA standards and check standards for each cation’s concentration (in parts per million)
(1,8). Plotting the values (Figure 1 and 2) provided a linear trend line that could be used to
determine the concentration values for both cations in parts per million. By using a tin foil piece
and a hot plate TDS (total dissolved solids) could be found. By using the “good transfer
technique” the samples, distilled water, and Ca2+ solution were placed on the tin foil until fully
evaporated. Samples were compared with distilled water and a Ca2+ (1.0x10^-3M) solution and
the results were recorded. The residue left behind determined the hardness of the water (the
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more residue, the harder the water). In some cases sample had to be diluted to proceed with
EDTA titration. EDTA titration involved placing different volumes of EDTA in each of a twelve
well set in order to examine color, and then determine concentration of the sample (as stated
above). Determining the concentration utilized the equation (1,8):
Equation 1: MEDTA x VEDTA = MCa2+ x VCa2+
This equation can be used to determine the concentration of Ca2+ in a sample using the number
of drops of EDTA taken to produce a color change in the well set. A similar equation was used
for Mg2+. Once the concentration was obtained, it was then converted to parts per million (ppm)
and grains per gallon (gpg) using the equations (1,8):
Equation 2: (M/V)*(100g/1 mole)*(1000mg/1g)=ppm
Equation 3: ppm*(1 grain per gallon/17.1 ppm)=grains per gallon
Once the color changed to blue in the well set, the determined the volume (measured in drops by
a previously calibrated microburet) of EDTA required to make the color change was plugged
into the equation 1 to determine concentration of the sample. EDTA titration was repeated to
note the change in hardness after the application of a commercial softening agent (baking soda)
and divalent cation exchange resin (1,8).
Results: (6,7,8,9)
The results of TDS for the four samples and control samples (distilled water and CaCl2) are listed
below. A description of each is provided in comparison to a reference sample.
Table 1.
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Total Dissolved Solids (TDS) Observations of Different Samples
Water Sample
Observation
CaCl2 Sample (Reference)
A feint white residue circle was left behind.
Distilled Water
Little to no white residue was left, less than
the reference.
Water taken from a tap on first floor of the
HUB (6)
A white ring of residue was left that was
darker than the reference with a defined outer
ring.
Water taken from a tap on second floor of
the HUB (7)
A white ring of residue was left that was
darker than the reference with a defined outer
ring.
Water taken from a tap on the third floor of
the HUB (8)
A white ring of residue was left that was
darker than the reference with a defined outer
ring.
Water taken from a filtration station on the
second floor of the HUB (9)
A white ring of residue was left that was
slightly darker than the reference, no defined
outer ring.
The calibration graphs of AA for Ca2+ (figure 1) and Mg2+ (figure 2) are shown below (8). The
graphs show absorbance values for calcium and magnesium (respectively) vs. concentration in
ppm. A trend line is shown with an equation for it, along with an R2 value to show point
deviation.
Figure 1.
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Light Absorbance vs. Ca+ ion Concentration
0.6
Light Absorbance
0.5
y = 0.0094x + 0.0079
R² = 0.9979
0.4
0.3
0.2
0.1
0
0
10
20
30
40
50
60
Ion Concentration (ppm)
Figure 2.
Light Absorbance vs. Mg+ ion Concentration
0.6
Light Absorbance
0.5
y = 0.0155x + 0.0162
R² = 0.9932
0.4
0.3
0.2
0.1
0
0
5
10
15
20
25
30
35
Ion Concentration (ppm)
Below is a table of absorbance values obtained through AA and the corresponding calculated
concentrations.
Table 2.
AA and Hardness Values Calculated for all Samples
Sample/Ions
AA Absorbance
Concentration
(ppm)
Hardness (ppm CaCO3)
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First Floor HUB
Tap (6)
Total: 230.8155
Mg2+
0.4996
31.1871
128.3420
Ca2+
0.3932
40.9894
102.4735
Second Floor HUB
Tap (7)
Total: 229.2492
Mg2+
0.4931
30.7677
126.6162
Ca2+
0.3938
41.0532
102.6330
Third Floor HUB
Tap (8)
Total: 241.6594
Mg2+
0.5149
32.1742
132.4041
Ca2+
0.4187
43.7021
109.2553
Second Floor
Filtration Station (9)
Total: 256.4138
Mg2+
0.5288
33.0710
136.0945
Ca2+
0.4603
48.1277
120.3193
To find the concentrations of Mg2+ and Ca2+ in parts per million the absorbance values were
plugged in as the y value in the equation on the graph that corresponds to the ion (8).
Equation 4: y=mx+b
Third Floor HUB sample calculations:
Mg2+ 0.5149=0.0155x+0.0162 x=32.174ppm
Ca2+
0.4187=0.0094x+0.0079 x=43.7021ppm
To find the overall hardness of the sample in terms of divalent cations the ppm values for Ca2+
and Mg2+ must be converted into ppm of CaCO3 through the equations shown below. The
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values are then added together to get the overall hardness in terms of ppm of CaCO3. The
calculations for the third floor tap of the HUB are shown below.
Equation 5: (ppm Ca2+)*[(100g CaCO3/1 mole)/(40.0g Ca2+/1 mole)]
(43.7021 ppm Ca2+)*[(100g CaCO3/ 1 mole)/(40.0g Ca2+/1 mole)]=109.2553 ppm CaCO3
Equation 6: (ppm Mg2+)*[(100g CaCO3/1 mole)/(24.3g Mg2+/1 mole)]
(32.1742 ppm Mg2+)*[(100g CaCO3/1 mole)/(24.3g Mg2+/1 mole)]=132.4041 ppm CaCO3
Equation 7: (ppm value 1)+(ppm value 2)=total hardness
(109.2553 ppm CaCO3)+(132.4041 ppm CaCO3)= 241.6594 ppm hardness
Table 3.
EDTA Results for Water Samples
Sample
Dilution Factor
(parts sample:
parts distilled
water)
Drops Needed for
Color Change
Hardness (ppm)
First Floor HUB
Tap (6)
1:3
5
400
Second Floor HUB
Tap (7)
1:1
7
280
Third Floor HUB
Tap (8)
1:3
4
320
Second Floor
1:1
Filtration Station (9)
9
360
To calculate the hardness CaCO3 from drops of EDTA the following equation below (Equation
8) was used. The calculations for the Third Floor of the HUB tap are shown (8).
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Equation 8: MEDTA * VEDTA = MSample * VSample
(2x10^-4M)*(4 drops EDTA)=MSample*(1 drop sample) Msample=(8.0x10^-4 M)*4 (dilution
factor)=3.2x10^-3M
(.0032 moles CaCO3/1 L solution)*(100.0g CaCO3/ mol CaCO3)*(1000 mg/g) =320ppm
Table 4.
The Hardness Values For Water Samples After Softening Techniques
Sample
Commercial Softening
Agent (ppm)
Divalent Cation
Exchange Resin (ppm)
First Floor Tap (6)
480
<96
Second Floor Tap (7)
576
192
Third Floor Tap (8)
288
80
Filtration Station (9)
384
200
The two water softening techniques were applied to all samples. The concentrations were
determined through the same EDTA titration equations above.
Discussion:
The results are slightly varied from technique to technique. The TDS test showed that the
taps of the three floors of the HUB had about the same residue. They all left a white circle with a
defined outer ring that was darker than the reference (CaCl2), meaning more impurities (cations)
were in the water samples. The second floor filtration station left a circle darker than the
reference, but not quite as dark as the other three samples. It makes sense that the filtered water
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would leave behind less residue. In the AA tests however, the filtration sample was the hardest
sample at approximately 256 ppm. The second hardest was the third floor tap sample at
approximately 242 ppm. The first and second floor taps showed extremely close results, both
near 230 ppm. Based on the AA tests, the filtration station seemed to produce water that was
surprisingly harder than the taps of the three floors in the HUB. The EDTA results followed the
same pattern as the AA analysis with the exception of the first floor tap. The first floor tap came
out much harder than the rest at 400 ppm. The next three were the filtration at 360 ppm,
followed by the third floor tap at 320 ppm, and finally the second floor tap at 280 ppm. While
these results show a steady decline of 40 ppm, the arrangement is quite random.
The hypothesis that the tap water samples from the three floors are equal in hardness was
incorrect. Also, it was incorrect in saying that the second floor filtration station would give a
softer sample of water. Based on the AA results in Table 2, the filtration station provided water
that was harder than the other three floors. With the exception of the first floor, the filtration
station was harder than the taps in EDTA as well. In the EDTA titration there may have been
sources of error that skewed the results of the first floor tap sample. The third floor appeared to
be harder than the other two taps. There may have been scale buildup in the upper floor’s piping
that lead to a harder water sample. If the pipes had been more recently replaced on the lower
floors the samples would be softer, as seen in the AA results. While it was assumed that the
filtration station would remove these divalent cations from the water, it proved false. The
filtration station may only remove large debris and particles and play no role on the
concentration of Mg2+ and Ca2+ along with other polyvalent cations in the water. Based on the
results it seems that the station increased the hardness of the water. An explanation for why the
samples gave these results is that the samples were obtained on different days of the week.
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Though the water comes from the same source, the water may have been harder on the day the
filtration station sample was taken compared to the other three. It may have also been harder on
the day the third floor tap water was taken in comparison to the first and second floor taps. All
of the samples had calcium hardness lower than the 2008 State College study, where the
hardness was between 180 and 220 ppm (10). Perhaps in the past few years measures have been
taken to soften the water in State College. The samples are still considered “hard” based on the
120-180 ppm scale (11).
EDTA can provide us with an analysis of total divalent cation concentration while AA
analyzes specific ions. AA seems to be more precise and accurate. EDTA has more room for
error because there is more human involvement and steps taken. Some possible sources of error
include the wells not being completely clean or dried. The wells could dilute the solutions
further where it may not be appropriate if they are not dried fully. They could also provide
unwanted contaminations if not properly cleaned. Another source of error is in the microburets,
considering not all drops come out the exact same size. Small pressure differences can produce
different sized drops, along with the speed at which the drops come out. Observational error
could play in when judging the color change from red to blue in the wells. People may observe a
well as blue when it is actually in an intermediate phase of red and blue. Also, if the appropriate
amount of EDTA to fully chelate the ions lies between two drops, the results are either too high
or too low. A source of error in the AA analysis is in the check standards. The AA values,
although close, were not exactly what they were supposed to be and so the absorbance values
may have been slightly inaccurate.
The two softening techniques worked in opposite ways. The commercial softening agent
actually increased the hardness of all four samples. There was perhaps error involved that
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produces inaccurate results. The agent may have been applied incorrectly, or it may have been a
faulty batch of product. The more effective of the two techniques was clearly the exchange
resin. The resin worked extremely well, softening the water by more than 100 ppm in two of the
samples (third and first floor taps), and more than 50 ppm in the other two (filtration station and
second floor tap). The resin, which is made primarily for substance exchange, worked much
better as it replaced divalent cations from the water with monovalent (softer) cations.
Conclusion:
Our hypothesis was almost entirely incorrect since the filtration proved to be the hardest
sample and the values for the taps of the three floors were not all the same. The only accuracy
was that the first and second floor taps had hardness values that were within 2 ppm of eachother.
The second floor filtration station had the hardest water. The third floor tap had the second
hardest water sample. The first floor was the third hardest by AA analysis, although the EDTA
titration showed it as being the hardest which is likely attributed to error. The second floor tap,
although very close to the first floor tap, was the softest sample. The results show that State
College water is relatively hard. The softening results showed that certain softening agents (in
this case cation exchange resin) can work remarkably to soften a water sample.
References:
(1)PSU Chemtrek by Stephen Thompson, Prentice Hall, Englewood Cliffs, NJ 2011-2012
version 10.0-10.22
(2)Briggs, J.C., and Ficke, J.F., 1977, Quality of Rivers of the United States, 1975 Water Year -Based on the National Stream Quality Accounting Network (NASQAN): U.S. Geological Survey
Open-File Report 78-200, page 436
(3) Water Hardness and Alkalinity http://water.usgs.gov/owq/hardness-alkalinity.html#briggs.
(accessed March 26, 2012).
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(4)Hard Water Problems http://www.hardwater.org/hard_water_problems.html (accessed March
26, 2012)
(5) Brown, Theodore. Chemistry: The Central Science, 11th Edition; Prentice Hall. Upper Saddle
River, 2009, pg 786
(6) Sam, Kyle, Chem 111 Laboratory Notebook, pp. 24-26
(7) Shaw, Mike, Chem 111 Laboratory Notebook, pp. 26-31
(8) Sampson, Dillon, Chem 111 Laboratory Notebook, pp. 28-32
(9) Sherman, Sydney, Chem 111 Laboratory Notebook, pp. 27-30
(10) State College Water may Cause Kidney Stones. http://voicesweb.org/node/1670. (July 7,
2008)
(11) Water Resource Characterization Dss. (n.d.). Hardness. (accessed March 27, 2012)
www.water.ncsu.edu/watershedss/info/hardness.html
(12) Facts About Your Water. City of Philadelphia. http://www.phila.gov/water/fact_sheets.html
(accessed March 29, 2012)