EXPERIMENT 3.3
WATER HARDNESS TITRATION
INTRODUCTION
One of the important contrasts between the alkali metals and alkaline earth metals
is the difference in the degree of solubility of the metal ions with respect to a variety
of anions. Given a metal from each group on the same row of the periodic table, the
higher charge density on the alkaline earth metal cation compared to the alkali
metal cation tends to allow the former to make stronger bonds with anions. This is
illustrated by the observation that nearly all ionic compounds containing alkali
metals are highly soluble in water whereas many alkaline earth salts are insoluble
or only slightly soluble.
Most sources of “fresh” water contain some concentration of alkaline earth cations,
mostly Ca2+ and Mg2+, that come from the slight solubility of these ions in the rocks,
soils, and minerals that are in contact with the water. It is a particular nuisance
that with variations in concentrations or in pH these ions may form precipitates
with anions already present in the water (such as hydroxide and carbonate) or ions
added during use of the water (such as phosphates or the fatty acid anions in
soaps). Bathtub ring and scale in water system pipes are examples of this type of
precipitation.
Water with a high concentration of Ca2+ and /or Mg2+ is called “hard” due to its
potential to produce solid precipitates. Water with low concentrations is called
“soft” water. A common measure of water hardness is obtained by determining the
combined concentration of these metal ions and imagining that it is all Ca2+ from
dissolved CaCO3. The concentration is expressed in milligrams of dissolved CaCO3
per liter of water sample. Since one liter of water has a mass of about one kilogram
and a milligram is one millionth of a kilogram, concentrations in milligram per liter
of water are often referred to as parts-per-million or ppm. Water hardness can vary
considerably from one water system to another. Values from tens to hundreds of
ppm are common in the continental United States.
In this experiment, you will measure the hardness of a sample of water. In order to
make this measurement by a titration, we need a reagent that will ignore alkali
metal ions but react with Ca2+ and Mg2+. One such reagent is EDTA (Ethylene
Diamine Tetraacetic Acid). Its structure will be described in a text chapter on
transition metals. For now, consider that it is a tetraprotic acid and that in a basic
solution, its anions tend to bond to many metal cations (but not alkali metal ions) in
a one-to-one ratio. Metal ions bound to EDTA are less available to form precipitates
with other anions so EDTA is used as a water softener in many shampoos and in
other applications.
We will use EDTA to measure the concentration of Ca2+ and Mg2+ (and any other
strong bonding cations) by adding it to water samples from a buret until the metal
ions are all bound. As an indicator, we will use a dye known as Eriochrome Black
T. Like EDTA it tends to bond to metals, however it bonds less strongly and to
fewer of them (e.g., Mg2+ but not Ca2+). Eriochrome Black T is a deep blue green by
itself. If it is added to a sample containing at least some Mg2+, it turns wine red as
they become bound together. This will be the initial sample color during a titration.
As EDTA is slowly added to a sample, it will begin tying up the metal ions. As long
as there are excess metal ions, the indicator can find them and assume the wine red
color. As the EDTA amount becomes equal to the metal ion amount, it will displace
the indicator from the last of the Mg2+. The indicator will then return to its blue
green color. This marks the endpoint.
As mentioned earlier, the water hardness measure treats the bonding metal ion
content as if it were all Ca2+. So we can treat the fundamental relationship in the
titration as
moles EDTA = moles Ca2+
Calculation of the moles of EDTA requires knowledge of its volume and molar
concentra-tion. Volumes will be measured with a buret. The concentration will be
found by titration against a standard solution of known Ca2+ equivalent
concentration that will be provided.
MATERIALS AND SAFETY
The standard and unknown solutions used do not contain corrosive substances.
However, the pH 10 buffer is considerably basic and constitutes an eye hazard.
Wear safety goggles at all times as usual. The Eriochrome Black T indicator is not
particularly toxic, but it will tend to stain. Handle it with care and clean up spills
promptly. All solutions used and produced in the experiment can be disposed of in
the sinks when finished.
PROCEDURE
PART A: STANDARDIZATION
Obtain a buret and a buret clamp from the equipment cart. Treat the burets with
great care, especially the fragile tips. Rinse the buret with distilled water, then
secure it to a vertical stand with the clamp. Assign one of your beakers to collect
rinsings. Use distilled water to verify that the stopcock functions properly with no
leaks or air bubbles.
Use a small beaker to transfer EDTA solution from the reservoir to the buret. Take
EDTA only as you need it to avoid wasting excess. Rinse the buret with very small
portions of EDTA before filling. To find the molarity of the EDTA it will be used to
titrate a known concentration of CaCO3. Use a 50 mL beaker to obtain about 40 mL
of the CaCO3 standard solution from the cart. Use a transfer pipet to deliver 10.00
mL of the standard into a clean 125 mL or 250 mL Erlenmeyer flask. Use a
graduated pipet to add 5 mL of pH 10 buffer (The instructor may have set out a
beaker and pipet for the entire class to use with the buffer). Next add 3-5 drops of
the Eriochrome Black T indicator. The standard solution has been laced with Mg2+
so the solution should appear wine red.
Record the initial liquid level in the buret. All readings should be recorded to the
nearest 0.01 mL. Add EDTA to the flask until the indicator changes from wine red
to blue green. Some may find this endpoint difficult to see at first. On the verge of
the endpoint the color is a somewhat complex mixture. Focus on the elimination of
traces of redness in it. It may help to compare the color to an untitrated (next)
sample and to saved, previously titrated samples. Record the final buret volume.
Repeat the process for a second and third trial. Refill the buret between trials.
Each partner should do at least one standard titration.
Try to perform operations consistently to enhance precision but do not be overly
concerned if it is not high. Several factors including the low concentrations and
interference from the buffer work against obtaining a sharp endpoint.
PART B: UNKNOWN ANALYSIS
In this experiment, you are usually given the option to use your own untreated
water sample or to draw from a reservoir of lab tap water. Indicate the source of
your sample (e.g., what city or water district or “lab water”) in the report section.
Due to the range of possible metal ion concentrations in the samples, choosing a
good sample size may require a trial-and-error approach. Ideally, a sample would
require 20-40 mL of EDTA to titrate. Start with a 200 or 250 mL sample for the
first trial, then adjust the sample size if necessary for the remaining trials. Be sure
to record the sample volume you use in each case. Before taking a sample, be sure
that the source is well-mixed by shaking or swirling.
Your instructor will indicate whether to use repeated transfers with a 50 mL pipet
or with your 50 mL graduated cylinder to deliver the sample volume into your 500
mL Erlenmeyer flask. Add buffer and indicator as with the standard. If at this
stage, the sample solution does not give the wine red color, consult your instructor.
You will either be directed to choose an alternative sample source or be instructed
in a method to add an Mg2+ spike and perform a spike-plus-blank titration to correct
for it. Assuming your sample did become wine red, titrate with EDTA as with the
standard recording the buret levels. If your group has one sample source, perform
at least three titrations with it. If two partners wish to analyze samples from two
different sources, perform two titrations on each. Then each partner should
determine the average Ca2+ molarity for their source only and record it under item
(o) of the calculations section of the report.
When finished, dispose of all solutions in the sink and rinse all glassware.
Carefully return the burets and pipets to the cart along with the buret clamps.
PRELAB QUESTIONS
1. A 10.00 mL aliquot of 0.010 M CaCO3 is titrated with 18.22 mL of EDTA
solution. What is the EDTA molarity?
2. A 250.0 mL water sample requires 30.85 mL of the EDTA solution from question
1 to reach the Eriochrome Back T endpoint. What was the molarity of the hard
metal ions in the water sample?
3. If the metal ions in the water sample of question 2 are assumed to be Ca2+ from
CaCO3, express the concentration in ppm CaCO3. This is the same as the
concentration in units of mg CaCO3 per liter of sample.
PART A (Create the following blank Tables in your lab notebook as part of prelab)
DATA AND POSTLAB CALCULATIONS
a) Molarity of standard Ca2+ solution______________
TRIAL
b) Initial EDTA level
1
2
3
___________ ___________ ___________
c) Final EDTA level
___________ ___________ ___________
d) EDTA volume delivered
___________ ___________ ___________
e) Qualitative comments
TRIAL
f) EDTA molarity
1
2
3
___________ ___________ ___________
g) Average EDTA molarity ± average deviation ___________ ± ___________
PART B (Create the following blank Tables in your lab notebook as part of prelab)
DATA AND POSTLAB CALCULATIONS
h) Source of water sample____________________________________
i)
TRIAL
Water sample volume
1
2
3
4
___________ ___________ ___________ ___________
j) Initial EDTA level
___________ ___________ ___________ ___________
k) Final EDTA level
___________ ___________ ___________ ___________
l) EDTA volume delivered
___________ ___________ ___________ ___________
(Corrected EDTA volume)
___________ ___________ ___________ ___________
m) Qualitative comments
TRIAL
n) Sample molarity as Ca2+
1
2
3
4
___________ ___________ ___________ ___________
o) Average Ca2+ molarity ± average deviation ___________ ±
___________
CONCLUSION
Assume sample water “hardness” was all Ca2+ from CaCO3.
p) Average mass, in grams, of CaCO3 per liter
___________ g/L.
q) Parts per million CaCO3 (i. e., mg CaCO3/L)
___________ ppm.