EXPERIMENT #11 Water Analysis: Alkalinitypeople.morrisville.edu/~habera/PDFFiles/CHEM122LabManual/11ALKALINITY2011.pdfIn this experiment the total alkalinity of a water sample will

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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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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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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



_____________

_____________

_____________

_____________

Initial pH of sample

_____________

_____________

_____________

_____________

Final pH of sample

_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________

Average alkalinity

_____________

Standard deviation

_____________


_____________

Calculations

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Part II: Colorimetric Titration

Determinations



_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________


_____________

_____________

_____________

_____________

Average alkalinity

_____________

Standard deviation

_____________


_____________

Calculations

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III. Titration Data

Buret Reading Volume Acid

Added

pH Buret Reading Volume Acid

Added

pH

____________ ____________ ____________ ____________ ____________ ____________

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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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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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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

Documents

EXPERIMENT #11 Water Analysis: Alkalinitypeople.morrisville.edu/~habera/PDFFiles/CHEM122LabManual/11ALKALINITY2011.pdfIn this experiment the total alkalinity of a water sample will