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EXPERIMENT #11
Water Analysis: Alkalinity
OBJECTIVES:
To learn to operate a pH meter
To perform a potentiometric titration
To perform a colorimetric titration
To plot a titration curve for an alkalinity determination
To determine the alkalinity of a water sample
Water, water, every where,
And all the boards did shrink;
Water, water, every where,
Nor any drop to drink.
Samuel Taylor Coleridge (1772-1834) from “The Rhyme of the Ancient Mariner”
BACKGROUND:
Water is perhaps the single most important substance on the face of the Earth.* Even the
oxygen of our atmosphere is made from water by the process of photosynthesis. The search for
extraterrestrial life includes a search for water. Water is an “industrial chemical” used on a scale
greater than any other industrial chemical. Water entering our homes and factories must meet
certain criteria; water exiting waste treatment plants must not harm the environment. Water quality
must be measured frequently, simply, and accurately.
A very good approximation of water quality can be made by three simple determinations:
alkalinity, hardness, and total dissolved solids. Alkalinity is the subject of this experiment; hardness
will be the subject of a future experiment.
Alkalinity is the ability of a water sample to react with acid; that is, its ability to neutralize
acid. Normally carbonate (CO32) and bicarbonate (HCO3
) are the acid reacting components of
water.
CO32 + H+ HCO3
(1)
HCO3 + H+ H2CO3 H2O + CO2 (2)
In the case of polluted samples, hydroxide (OH) may also contribute to a sample’s alkalinity. It is
the water chemist’s practice to report alkalinity in terms of its equivalent in calcium carbonate
(CaCO3). The unit of alkalinity is mg CaCO3 per liter or simply mg/L. For dilute aqueous solutions
1 mg/L is the same as 1 ppm (part per million). Alkalinity is important in the treatment and
purification of domestic and industrial water and wastewater.
Water chemists also find it convenient to work with units known as equivalent mass,
equivalents, and normality. A simplified definition of equivalent mass for our purposes is the mass
equal to one mole of positive (or negative charge). For example: the equivalent mass of Al3+ is 9.0
g/eq (the molar mass 27.0 g/mol divided by the magnitude of the charge, 3); the equivalent mass of
CO32 is 30.0 g/eq; the equivalent mass of sulfuric acid (H2SO4, molar mass 98.0 g/mol) is 49.9
g/eq. Equivalents equals mass divided by equivalent mass. Normality is equivalents divided by
liters of solution. Note the analogy between equivalents/moles and normality/molarity.
*Philip Ball, Life's Matrix : A Biography of Water, June 2000, Farrar Straus & Giroux
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EXPERIMENT #11 WATER ANALYSIS: ALKALINITY
In this experiment the total alkalinity of a water sample will be determined by titrating the
sample with a known solution of sulfuric acid to a pH value of 4.50. When using a pH meter, the
analysis is known as a potentiometric titration; when using an indicator it is known as a colorimetric
titration. The volume of titrant used to reach the end point, the acid concentration, and the sample
volume are used to calculate the alkalinity as shown in the following equation.
alkalinity,mg LCaCOmL titrant normalityof acid 50,000
mLsample3
(3)
The relationship among the three contributors to alkalinity is shown below.
Hydroxide
Carbonate carbonate carbonate phenolphthalein end point pH 8.3
bicarbonate bicarbonate
pH 4.5
(a) (b) (c) (d)
Graphical representations of various forms of alkalinity and titration end point relationships*
If P equals the measured phenolphthalein alkalinity and T equals the total alkalinity, then
(a) hydroxide = 2P – T and carbonate = 2(T – P) (b) carbonate = 2P = T
(c) carbonate = 2P and bicarbonate = T – 2P (d) bicarbonate = T
EXPERIMENTAL PROCEDURE: (Work individually.)
Part I
Potentiometric Titration
1. Following the directions of your instructor calibrate a pH meter using the pH 7 and pH 4
buffers. Leave the pH electrode in one of the buffer solutions or deionized water.
NOTES 1. The filling hole of the electrode must be in the open position for proper use.
2. Keep the pH electrode immersed in distilled water or storage buffer (pH 7) whenever it is not being used.
3. Do not let the electrode dry out.
*Modified from Water and Wastewater Technology, M.J. Hammer and M.J. Hammer, Jr., Prentice Hall
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EXPERIMENT #11 WATER ANALYSIS: ALKALINITY
2. Set up your buret to properly deliver the titrant. The buret must be clean and without obstruc-
tions or leaks. Rinse the buret three times with the solution it will dispense. Fill the buret and
remove any air bubbles. Record the concentration of titrant and the initial volume.
3. Manipulate the buret, buret clamp, stand, pH meter, pH electrode, magnetic stirrer, magnetic
stir bar, and beaker into a stable configuration. See figure below.
4. Obtain a water sample. Record the sample or unknown number on your data sheet. Add 100
mL of sample to the beaker.
5. Start the stirrer. Measure the initial pH of the sample and record the value on your data sheet.
Switch the pH meter to the continuous mode and press the pH button. Leave the pH electrode
immersed in the beaker and continue stirring gently for the duration of the titration. Add titrant
until a pH of 4.50 is displayed on the pH meter. Be patient towards the end of the titration,
since the pH changes very rapidly at this stage and it is very easy to overshoot the end point.
Calculate the alkalinity using equation (3).
6. Rinse the beaker, the stir bar, and the pH electrode with deionized water. Blot the electrode
dry, and then immerse it into another 100 mL of the same sample. Place the stir bar into the
beaker and turn on the stirrer. Repeat step 5 as many times as it takes to get consistent results.
Refill the buret as necessary. Calculate the alkalinity for each trial, the average, the standard
deviation (sd), and the rsd (relative standard deviation). Your calculator is programmed to
perform these operations (see page 116). Ask your instructor for assistance.
7. When you complete your potentiometric titrations, promptly remove the pH electrode from the
solution, rinse it thoroughly with deionized water, and place in the pH 4 or pH 7 buffer.
Read the bottom of the
meniscus at eye level.
Add measured sample to a 250 mL
beaker.
Titrate to pH 4.5.
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EXPERIMENT #11 WATER ANALYSIS: ALKALINITY
Part II
Colorimetric Titration
1. Set up your buret to properly deliver the titrant. The buret must be clean and without obstruc-
tions or leaks. Rinse the buret three times with the solution it will dispense. Fill the buret and
remove any air bubbles. Record the concentration of titrant and the initial volume.
3. Manipulate the buret, buret clamp, stand, magnetic stirrer, magnetic stir bar, and beaker into a
stable configuration.
4. Obtain a water sample. Record the sample or unknown number on your data sheet. Add 100
mL of sample to the beaker. Add 3-5 drops of indicator -- a mixture of bromocresol green and
methyl red. The indicator changes from bluish-gray at pH 4.8 to pale yellow/light pink at pH
4.6
5. Start the stirrer. Add titrant until the proper color change is observed. Be patient towards the
end of the titration, since the pH changes very rapidly at this stage and it is very easy to
overshoot the end point. Calculate the alkalinity.
6. Rinse the beaker and the stir bar with deionized water. Add 100 mL of the same sample, place
the stir bar into the beaker and turn on the stirrer. Repeat step 5 as many times as it takes to get
consistent results. Refill the buret as necessary.
Part III (Optional)
Titration Curve
1. Set apparatus and sample as in Part I. Record the initial pH, then switch to continuous mode.
2. Add 1.00 mL of acid to your water sample, and record the pH of the solution when the readout
of the pH meter stabilizes.
3. Continue adding 1.00 mL increments of acid until within 2 mL of the pH 4.5 endpoint. After
each addition record the buret reading, the total volume of acid added, and the pH.
4. When you are within 2 mL of the equivalence point, add the acid in 0.20 mL increments.
Again, after each addition, record the buret reading, the total volume of acid added, and the pH.
5. When you complete your titrations, promptly remove the pH electrode from the solution, rinse
it thoroughly with deionized water, and place in the pH 4 or pH 7 buffer.
6. On the graph paper provided at the back of your lab manual, construct a plot of pH (vertical
axis) versus volume of acid added (horizontal axis).
Properly Reading a Buret
1. See the figure on page 130 for assistance in properly reading a buret.
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NAME________________________________________ Section_______ Date__________
DATA AND CALCULATIONS: Water Analysis: Alkalinity
Part I: Potentiometric Titration
Acid used as titrant: ____________ Sample identification: ____________
Concentration of standard acid solution: ____________
Practice Determinations
First Second Third Fourth
Volume of sample, mL
_____________
_____________
_____________
_____________
Initial pH of sample
_____________
_____________
_____________
_____________
Final pH of sample
_____________
_____________
_____________
_____________
Final volume of acid, mL
_____________
_____________
_____________
_____________
Initial volume of acid, mL
_____________
_____________
_____________
_____________
Volume of acid used, mL
_____________
_____________
_____________
_____________
Alkalinity, mg/L CaCO3
_____________
_____________
_____________
_____________
Average alkalinity
_____________
Standard deviation
_____________
Relative standard deviation
_____________
Calculations
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Sample identification: ____________
Unknown Determinations
First Second Third Fourth
Volume of sample, mL
_____________
_____________
_____________
_____________
Initial pH of sample
_____________
_____________
_____________
_____________
Final pH of sample
_____________
_____________
_____________
_____________
Final volume of acid, mL
_____________
_____________
_____________
_____________
Initial volume of acid, mL
_____________
_____________
_____________
_____________
Volume of acid used, mL
_____________
_____________
_____________
_____________
Alkalinity, mg/L CaCO3
_____________
_____________
_____________
_____________
Average alkalinity
_____________
Standard deviation
_____________
Relative standard deviation
_____________
Calculations
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NAME________________________________________ Section_______ Date__________
Part II: Colorimetric Titration
Determinations
First Second Third Fourth
Volume of sample, mL
_____________
_____________
_____________
_____________
Final volume of acid, mL
_____________
_____________
_____________
_____________
Initial volume of acid, mL
_____________
_____________
_____________
_____________
Volume of acid used, mL
_____________
_____________
_____________
_____________
Alkalinity, mg/L CaCO3
_____________
_____________
_____________
_____________
Average alkalinity
_____________
Standard deviation
_____________
Relative standard deviation
_____________
Calculations
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NAME________________________________________ Section_______ Date__________
III. Titration Data
Buret Reading Volume Acid
Added
pH Buret Reading Volume Acid
Added
pH
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NAME________________________________________ Section_______ Date__________
ADDITIONAL ASSIGNMENT I: Water Analysis I: Alkalinity
1. Pick one of the URL’s listed below and follow the assignment.
http://www.nbif.org/education/post-sec/env-engr/alkalinity/alkalinity.html
INTRODUCTION TO ENVIRONMENTAL ENGINEERING MODULE I. NATURAL SYSTEMS
ASSIGNMENT: Print out and submit the table of contents of MODULE I: Natural Systems for the
above course.
http://www.epa.gov/owow/estuaries/monitor/chptr11.html
Volunteer Estuary Monitoring
A METHODS MANUAL
ASSIGNMENT: Print out and submit the table of contents of the above manual.
http://water.usgs.gov/owq/FieldManual/Chapter6/6.6_contents.html
National Field Manual
By D.B. Radtke, F.D. Wilde, J.V. Davis, and T.J. Popowski
ASSIGNMENT: Print out and submit the table of contents of the section titled --
6.6 ALKALINITY AND ACID NEUTRALIZING CAPACITY
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EXPERIMENT #10 WATER ANALYSIS: ALKALINITY
ERROR ANALYSIS:
It is rare that a chemist will perform an analysis consisting of one determination. Several
samples, making up what is called a replicate analysis, are subjected to the analytical procedure.
The chemist reports the average of the several determinations and a number reflecting the precision
-- the reproducibility of each determination -- of the analysis. A common means of reporting
precision is to calculate the standard deviation of the several determinations. This approach is best
illustrated by the following example.
The results of an analysis are 2.100, 2.110, and 2.105. First calculate the average (also
known as the mean) of the three determinations.
2.1053
2.1052.1102.100mean
Next, calculate the deviation of each trial from the mean:
|2.105 - 2.100 | = 0.005
|2.105 - 2.110 | = 0.005
|2.105 - 2.105 | = 0.000
The standard deviation is calculated as follows
2
(0.000)(0.005)(0.005)sd
222
= 0.005
(The numerator under the radical symbol is the sum of the squares of the deviations. The
denominator is one less than the number of replicate samples.) The results of the above analysis
would be reported as 2.105 ±0.005. The relative standard deviation, rsd, is
ppm2000orppt2or0.2%or0.002or0.002382.105
0.005rsd
The standard deviation may be used to determine whether one or more of the replicate analyses
should be retained or discarded. If any result is more than two standard deviations from the mean,
then it is discarded. In the above example, none of the analyses would be discarded.
It is important that you learn how to do this mathematical manipulation on your
calculator. The example given above is oversimplified to illustrate the mathematical operations.
There will be significant round-off error if you try to perform these operations with your
experimental data without properly using your calculator. Your instructor will be glad to show you
the proper method for performing a single variable analysis.
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NAME________________________________________ Section_______ Date__________
ADDITIONAL ASSIGNMENT II: Water Analysis I: Alkalinity
1. Complete the following table. Make proper use of significant figures.
Ka
pKa
Ka
pKa
3.6 x 10-6
4.17
1.37 x 10-4
4.85
1.00 x 10-5
5.41
1.48 x 10-4
7.20
2. Fill in the missing portions in the following table.
Name Symbol/Formula Molar Mass Valence/Charge Equivalent Mass
Aluminum
Chromium +3
Chromium +6
Hydrogen
Nitrogen -3
Oxygen
Selenium +6
Ammonium ion
Hydroxide ion
Carbonate
Bicarbonate
Phosphate
Dihydrogen
phosphate
Sulfate
Bisulfite
Nitrate
Hypochlorite
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2. The alkalinity of a water sample consists of 16 mg/L of CO32 and 120 mg/L of HCO3
.
Compute the alkalinity in milligrams per liter as CaCO3. (HINT: Convert to equivalents, then
back to mass.)
3. Calculate the molecular and equivalent masses of ferric sulfate [iron(III) sulfate].
4. Calculate the alkalinity of the following water samples:
(a) (b) (c) (d) (e)
normality of acid 0.0200 0.0200 0.1000 0.0200 0.150
sample size, mL 50.0 100.0 175.0 150.0 85.0
mL acid to pH 4.5 10.00 27.32 17.53 26.50 16.73
alkalinity, mg/L CaCO3
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