The Chemistry Of Natural Water

INTRODUCTION
 
The purpose of this experiment is to explore the hardness
of the water on campus. Hard water has been a problem for
hundreds of years. One of the earliest references to the
hardness or softness of water is in Hippocrates' discourse
on water quality in 5th Century B.C. Hard water causes many
problems in both the household and the industrial world. 

One of the largest problems with hard water is that it
tends to leave a residue when it evaporates. Aside from
being aesthetically unpleasant to look at, the build up of
hard water residue can result in the clogging of valves,
drains and pipes. This build up is merely the accumulation
of the minerals dissolved in natural water and is commonly
called scale. 

Other than clogging plumbing, the build up of scale poses a
large problem in the industrial world. Many things that are
heated are often cooled by water running through piping.
The build up of scale in these pipes can greatly reduce the
amount of heat the cooling unit can draw away from the
source causing a potentially dangerous situation. It also
greatly reduces the heat efficiency of the container which
holds the water, therefore much more energy is required to
heat the item to the necessary temperature. In the
industrial world, this could amount to large sums of money
being thrown into wasted heat. 

In addition to clogging plumbing and reducing heating
efficiency, the build up of hard water also adversely
affects the efficiency of many soaps and cleansers. The
reason for this is because hard water contains many
divalent or sometimes even polyvalent ions. These ions
react with the soap and although they do not form
precipitates, they prevent the soap from doing its job.
When the polyvalent ions react with the soap, they form an
insoluble soap scum. This is once again quite unpleasant to
look at and stains many surfaces.
 
The reason for all these problems is that hard water tends
to have higher than normal concentrations of minerals, and
hence it leaves a considerable amount more residue than
normal water. The concentration of these minerals is what
is known as the water's Total Dissolved Solids or TDS for
short. This is merely a way of expressing how many
particles are dissolved in water. 

The TDS vary from waters of different sources, however,
they are present in at least some quantity in all water,
unless it has been passed through a special distillation
filter. The relative TDS is easily measured by placing two
drops of water, one distilled and one experimental on a
hotplate and evaporating the two drops. You will notice
that the experimental drop will leave a white residue. This
can be compared to samples from other sources, and can be
used as a crude way of measuring the relative TDS of water
from a specific area. The more residue that is left behind,
the more dissolved solids were present in that particular
sample of water. 

Another, perhaps more quantitative way of determining
hardness of water is by calculating the actual
concentrations of divalent ions held in solution. This can
be done in two ways. One is by serially titrating the water
with increasing concentrations of indicator for Mg++ and
Ca++ (we will be using EDTA). This will tell us the
approximate concentration of all divalent ions. The method
of serial titrations is accurate to within 10 parts per
million (ppm). 

A second possible method for determining the hardness of
water is by using Atomic Absorption Spectrophotometry or AA
for short. AA is a method of determining the concentrations
of individual metallic ions dissolved in the water. This is
accomplished by sending small amounts of energy through the
water sample causing the electrons to assume excited
states. When the electrons drop back to their ground
states, they release a photon of energy which can be
measured by a machine and matched up to the corresponding
element with the same E as was released. When the finding
is related to the intensity of the light emitted and the
amount of light absorbed, a concentration value is
assigned. 

A quick overview of how the atomic absorption
spectrophotometer works follows:
 
First, the water sample is sucked up. Then the water sample
is atomized into a fine aerosol mist. This is in turn
sprayed into an extremely high intensity flame of 2300 C
which is attained by burning a precise mix of air and
acetylene. This mixture burns hot enough to atomize
everything in the solution, solvent and solute alike. A
light is emitted from a hollow cathode lamp. The light is
then absorbed by the atoms and an absorption spectrum is
obtained. This is matched with cataloged known values to
attain a reading on concentration. 

Because there are so many problems with hard water, we
decided that perhaps the water on Penn State's campus
should be examined. My partners and I decided to test
levels of divalent ions (specifically Mg++ and Ca++ ) in
successive floors of dormitories. We hypothesized that the
upper level dormitories would have lower concentrations of
these divalent ions because being heavy metals, they would
tend to settle out of solution. The Ca++ should settle out
first seeing how it is heavier than the Mg++, but they will
both decrease in concentration as they climb to higher
floors in the dormitories.
 
PROCEDURE
 
We collected samples from around Hamilton Halls and West
Halls. In order to be systematic, we collected samples in
the morning from the water fountains near the south end of
the halls. We also collected water samples from each floor
for comparison. The reason we collected them in the morning
was so that the Mg++ and Ca++ would be in noticeable
quantities. We then went about and tested and analyzed via
serial titrations and via Atomic Absorption
Spectrophotometry. We also obtained a TDS sample merely for
the sake of comparison, and to ensure that there were in
fact dissolved solids in our water samples (without which
this lab would become moot). 

For the serial titration, we merely mixed the water sample
with EBT, and then with increasing concentrations of EDTA.
The EBT served as an indicator to tell us when the
concentrations of the EDTA and the divalent ions in
solution were equal (actually it told us when Mg++ was
taken out of solution but that served the same purpose).
This allowed us to find the concentration of the divalent
ions dissolved in solution. 

Based on this, we calculated the parts per million and the
grains per gallon for each water sample. Finally, we took
an AA reading for each sample. This gave us absorption
values and concentration values for each of the two main
metals we were observing; Ca++ and Mg++. We then plotted a
graph of Atomic Absorption Standards using values that were
given to us by the AA operator. These values helped us to
calibrate the machine. The parts per million that we find
will be based on plugging in the reported absorption value
into the resulting curve from the graph of these values. 

The resulting concentration was used as the final value for
the hardness for that particular sample. All calculations
and conclusions were done based on these final values
obtained for the concentration of Ca++ and Mg++. 

RESULTS
 
Molarity x (100g CaCO3 / 1 mole CaCO3 ) x (1000 mg / 1g) =
Xmg/1000g = ppm
 
Grains/Gallon = ppm /17.1
 
Example:
 
(1.6 x 10 -3 moles / 1 Liter) x (100g CaCO3 / 1 mole CaCO3
) x (1000 mg / 1g) = 160 ppm
 
160 ppm/17.1 = 9.35 grains/gallon
 
Conversion Factors Given by AA operator: Ca++ = 2.5
 
Mg++ = 4.2 Ca++ x 2.5 = CaCO3 hardness ppm value Mg++ x 100
x 4.5 = Mg CO3 hardness ppm value *NOTE: the Mg++ is x 100
because it was diluted before it was put into the AA.
 
CONCLUSION
 
Upon completion of this lab, it can be said that the data
supports only half of the original hypothesis. Yes, the
Ca++ did seem to decrease as the water got further from the
source and climbed higher in the dormitories. However, the
Mg++ did not. Instead it did quite the opposite and showed
a general trend of increasing in concentration as it got
further away from the source and higher in the dormitories.
Perhaps a viable explanation could be attained if studies
were done on the plumbing inside the building. Perhaps
there is a high concentration of magnesium in the solder
used to hold the pipes together. Perhaps it is not in the
pipes but rather perhaps the people on the upper floors get
up later and therefore at the time of collection, the water
in the upper floors had been sitting longer than that on
the lower floors. 

In either case, more investigation would have to be
conducted in order determine what caused the unexpected
results. In light of this discrepancy, the overall accuracy
of the lab was very good. The numbers all seem to back each
other up and correlate very well. As was mentioned in the
previous section, the precision and accuracy with which
this lab was carried out seems to indicate that there is
very little source of error. Overall, it seems that the lab
was quite well done. The hypothesis would have to be
revised and as of this point, without further
investigation, it would have to be reformulated to say that
only the Ca++ would decrease in concentration whereas the
Mg++ would increase.
 
REFERENCES:
 
Brown, Theodore L. et al. Chemistry The Central Science;
Sixth Edition; Prentice Hall, Englewood Cliffs, NJ; ©1994