Determination of Water Hardness
Complexometric Titration with EDTA
Purpose:
You will prepare and standardize a disodium EDTA solution and use it to determine the
hardness of water samples.
Overview:
Water hardness is an expression for the sum of the dissolved calcium and magnesium
cation concentrations in water. These cations form insoluble salts with soap, known as
soap scum, decreasing soap’s cleaning effectiveness. Also, when hard water is heated,
calcium carbonate precipitates, clogging pipes and water heaters. Although it is a
potential nuisance, drinking hard water is not a health hazard. Drinking hard water
contributes a small amount of calcium and magnesium toward the total human dietary
needs. People regularly take calcium supplements to increase their calcium intake. The
standard way to express water hardness is in ppm (parts per million) as CaCO3.
An excellent way to determine water hardness is to perform a complexometric titration
using a standard ethylenediaminetetraacetic acid (EDTA) solution. Titration is the
volumetric measurement of a solution of
known concentration when it reacts
completely with a measured volume or
mass of another substance.
Complexometric titrations rely on a
complexation reaction. In complexation
reactions metal ions act as Lewis Acids
(electron pair acceptors), which bind
with Lewis bases (electron pair donors)
called ligands. If the ligand possesses
more than one donor atom it can be called a chelating
agent. The chelating agent used in this experiment is
the ethylenediaminetetraacetate ion [EDTA] 4 - . It has
six donor atoms.
Ethylenediamine can be though of as two ammonia
molecules bound together with a two atom carbon
chain. Each nitrogen has a lone pair of electrons so
each is capable of forming a coordinate covalent bond
with metal ions. Each nitrogen is also bound to two
acetate ions1, CH3CO2-. One of the oxygens on each
acetate is also capable of forming a bond to the metal
ion. Since there are four such acetate ions, we now
have six bonds to the metal ion. The EDTA literally
wraps around the metal ion and binds it quite strongly.
1
Note that the acetate is missing one hydrogen atom and is bound to the nitrogen instead.
In solution EDTA can pick up hydrogen atoms and exist as [EDTA2+]; however, this
species does not act as a chelating agent. Therefore, it is necessary to keep the pH of the
titration at 10 in order to drive the equilibrium of the compound to the [EDTA] 4-. A
buffer is used to maintain the desired pH.
In order to determine the endpoint in the titration a metal ion indicator is used. This
indicator will also form a complex with the metal ion, though not as strongly as the
[EDTA4-]. When EDTA is first added to the solution it will bind with the free metal ion,
but as the supply of free metal ion depletes it will begin to react with the metal ion that is
bound to the indicator.
Metal*Indicator + EDTA → Metal*EDTA + Indicator
In this experiment the indicator that is used is calmagite2; which, when bound to calcium
forms a pink complex, and when not bound to calcium forms a blue complex. Due to
steric hindrances3, EDTA will complex with calcium and magnesium in a one-to-one
molar ratio.
Calmagite
In the first part of this experiment you will prepare and standardize an EDTA solution.
You will then use the standardized EDTA solution to measure the hardness of water
samples.
Materials:
50 mL graduated cylinders
500 mL and 1000 mL volumetric flasks
50 mL buret
25 mL pipette
Erlenmeyer flask
funnel
ethylenediaminetetraacetic acid, disodium salt dihydrate
calcium carbonate
concentrated hydrochloric acid
calmagite indicator
water samples
2
1-(1-hydroxy-4-methyl-2-phenylazo)-2-naphthol-4-sulfonic acid.
steric hindrance is physical blockage of a particular site within a molecule by the presence of local atoms
or groups of atoms
3
Safety Precautions:
Make sure that you wear safety goggles and apron at all times. Acids and bases are
corrosive chemicals and can cause burns to skin and eyes. Avoid contact and wash any
contaminated area thoroughly with cold water. Report any spills.
Procedure:
Preparation of EDTA Solution
1)
Prepare an approximately 0.01 M solution of the disodium salt of EDTA
(molar mass 372.24 g/mol). First, add the appropriate amount of dry4
disodium EDTA to a 500 mL volumetric flask. Then, add deionized water to
the line and mix.
Standardization of EDTA Solution
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
11)
Add 1g (record mass to .001g if possible) of dry5 calcium carbonate to a
1000mL volumetric flask.
In the fume hood, add 2 mL of concentrated HCl to the flask to dissolve the
calcium carbonate.
Add deionized water to the line and mix. You now have an approximately
0.01M CaCl2 solution.
Using a funnel, rinse and then fill the buret with about 50mL of the EDTA
solution. Open the stopcock and allow a few mL to drain through the tip into
a beaker to flush out any trapped air bubbles. When the solution has settled in
the buret, record the volume to ±0.01mL reading the level from the top down.
(Remember significant figures: record numbers from the markings and a last
number estimated between the final two markings.)
Accurately measure 25 mL of the 0.01M calcium chloride solution into a
250mL Erlenmeyer flask. Record the exact volume.
Add 1.0mL of the ammonia buffer6 solution into the flask.
Add several drops of the calmagite indicator. The solution should turn a pink
color.
Record the initial volume of the EDTA solution in the buret.
Titrate until the solution turns a light blue color. It will turn purple before
becoming blue. Do not over titrate.
Record the final volume of the EDTA solution.
Calculate the concentration of the EDTA solution.
MolarityEDTA x VolumeEDTA = VolumeCa solution x MolarityCa solution
12)
4
Repeat the titration until results are consistent. Average your results.
dried at 80ºC for 1 week.
dried for at least 1 hour at 150ºC
6
pH 10 ammonia buffer solution prepared from 57.0mL concentrated ammonium hydroxide, and 6.75g
ammonium chloride in 100mL DI water.
5
Titration of Samples
1)
2)
3)
4)
5)
6)
7)
Accurately measure 25 mL of the water sample into a 250mL Erlenmeyer
flask. Record the exact volume.
Add 1.0mL of the ammonia buffer solution into the flask.
Add several drops of the calmagite indicator. The solution should turn a pink
color.
Record the initial volume of the EDTA solution in the buret.
Titrate until the solution turns a light blue color. It will turn purple before
becoming blue. Do not over titrate.
Record the final volume of the EDTA solution.
Calculate the concentration of the calcium in your solution.
MolarityEDTA x VolumeEDTA = VolumeCa solution x MolarityCa solution
8)
9)
Repeat the titration until results are consistent. Average your results.
Continue by repeating the titration for your other samples.
Interpretation
1) Convert your calcium ion concentrations to milligrams per liter of CaCO3. This is
approximately equal to ppm of CaCO3.
2) Using the table below, determine if the water samples you tested would be
considered soft, moderately hard, hard, or very hard.
Water Hardness Table7
(Concentrations as CaCO3)
In Grains per Gallon (gpg)
Below 3.5
3.5 to 7.0
7.0 to 10.5
10.5 and above
In Milligrams per Liter (mg/L)
Below 60
60 to 120
120 to 180
180 and above
Relative Hardness Level
Soft
Moderately Hard
Hard
Very Hard
Developed from Bowling Green State University, Pacific University, The University of Tennessee, College of Charleston laboratory experiments:
http://www.chem.pacificu.edu/Projects2005/Pages/CalciumTitration.htm
http://www.bgsu.edu/departments/chem/faculty/endres/ch128/Calcium.htm
http://www.chem.utk.edu/~chem319/Experiments/Exp6.pdf
http://www.cofc.edu/~kinard/221LCHEM/CHEM221L%20Calcium%20Determination%20by%20EDTA%20Titration.htm
7
source: Virginia Cooperative Extension http://www.cdc.gov/nasd/docs/d001201-d001300/d001228/d001228.pdf
Questions:
1) How hard are your water samples?
2) At pH 10 calcium has a +2 charge and EDTA ion has a -4 charge. How is the
stoichiometric ratio 1:1 for EDTA:Ca2+?
3) How do you think that the calmagite bonds with the calcium?
4) What is the purpose of the buffer?
5) What would happen if no buffer was added?
6) How would it have affected your results, if the EDTA had not been completely
dry?
7) How would your results be affected if the deionized water was contaminated with
calcium, magnesium, or other metal ions?